So far you've learned that atoms are composed of subatomic particles such as protons, neutrons, and electrons. In this lecture I'll introduce you to the concept of ions, species in which the number of electrons does not equal the number of protons. We'll also talk about isotopes, species that differ by, the number of neutrons that are present. Now, atoms that have a charge, that is not zero, are called ions. Recall from previous lectures, that the number of protons determines the identity of the element. So for example, all atoms of sodium have 11 protons in the nucleus. And if you look on the periodic table, the atomic number for sodium, the symbol is Na, remember, is eleven. Your probably going to need a periodic table to follow this lecture and do some of the in-video excersises. So, go ahead and pause the video to get out a periodic table if you need to do so. Atoms are not neutral and bear an overall net charge if the number of protons does not equal the number of electrons. So, for example, in the last lecture, I noted, I noted that the alkali metals, such as lithium, sodium, and potassium, tend to lose one electron. They give those electrons up fairly easily. An atom bears an overall net charge, if the number of protons does not equal the number of electrons. For example, if a sodium atom has only 10 electrons, and not 11, then its charge is +1. And because it has a positive charge, more protons than electrons, it is called a cation. The cation is a species that has a positive charge. There are some important points about ions that were repeating. The first thing is that when a species forms an ion, it does so by gaining or losing electrons, not protons. If the number of protons is changing, then the identity of the element would change. But the number of electrons can change and the identity of the element remains the same, it just changes into an ion. The ions charge can be easily calculated. It's the number of protons in the species minus the number of electrons. Again, any time an atom has a positive charge its called a cation. Sometimes its a collection of atoms that all together have a positive charge. And we still call that a cation. A cation is any species where there's more protons than electrons. Cations are formed by the species losing electrons at some point in the past. If an ion has a negative charge, then it's called an anion. An anion always has more electrons than protons. It could be a monatomic anion, meaning there's a single atom that has more electrons than protons, or it could be a collection of atoms where the total number of electrons is greater than the total number of protons, and that's called a polyatomic ion. Again, anions form whenever a neutral atom gains electrons. Non-metals tend to be more attractive to electrons than metals. So, non-metals tend to form anions more readily than metals do. Let's do some review and practice. Again, you're going to probably want to have a periodic table out for this exercise. You can find a periodic table on the main page of the Coursera website or typically it's on the cover of any general chemistry textbook. Let's start with a simple question. On your periodic table look and find which element has 14 protons. Great, I see you've taken time to find the periodic table and you found that the element with the atomic number of 14 is Silicon. That one wasn't too bad, was it? Now let's do some more complicated examples, where we look at the number not just of protons, but also of electrons and neutrons, in some atomic species. So, in this table, which we're going to fill out together, we're going to start by being given a species, and then we're going to determine the number of protons, electrons and neutrons than that, that that species has. For example, zinc, how many protons does zinc have? Well, it might take you a minute to find it on your periodic table, but go ahead and do that. Now that you found it, you see that zinc always has 30 protons. The number of protons is what makes a particular atom, an atom of zinc. Now, neutral zinc has the same number of electrons as it does protons. So, the negative charges from the electrons completely cancel out the positive charges from the protons. Therefore, if I've just written zinc as I have here with no charge shown, I have to assume that it's neutral zinc, and neutral zinc must have thirty electrons. How many neutrons does an average zinc atom have? Let's think about that. How would we figure that out? Well, we would look at the periodic table. Here's the box on the periodic table that's got zinc in it. At the top of the box is the atomic number, at least on my periodic table. The smaller number without the decimal place after it is always the atomic number, and that's where we got this 30 for the number of protons. The other number that's usually shown on a periodic table is, remember, the average atomic mass. Sometimes that is not a round number. For example, here it's 65.39. If we know that protons all have a mass of about 1, mass approximately equal to 1 and we know that neutrons all have a mass of approximately equal to 1 and electrons are so light that we don't bother to worry about them. When I say one, I should put a unit. Those are all both one atomic mass unit. Electrons are approximately 10 to the minus 4 atomic mass units each. So, we're just going to ignore the number of electrons. If we look at the mass of a zinc atom we can see it's 65.39 atomic mass units, which seems a little strange, so let's just round it down to 65. If the mass of zinc is 65 atomic mass units and we subtract the number of protons, which should weigh 30 atomic mass units, we see that there are probably 35 atomic mass units left, and those all come from the 35 neutrons that are present. So I can calculate the number of neutrons by taking the atomic mass and subtracting the atomic number See what I did? Because the electrons are light, and they don't really contribute very much to the mass of the atom. All right, that's a neutral atom and this lecture's supposed to be about ions. Let's do an example of an ion. O2 -, so this oxygen has a charge on it. How many protons does this oxygen contain? Well, if it's oxygen, we can look on the periodic table and see that the atomic number is 8. It needs to have eight protons, or it's not oxygen. How many electrons does this oxygen atom contain? Well if it had 8 electrons that would be neutral oxygen. This oxygen is not neutral. It has a minus 2 charge. Sometimes you'll see it written as -2. Sometimes you'll see it written as 2-. Those are both the same as far as I'm concerned. Therefore, this particular oxygen has more electrons than protons, two more, in fact, for each atom. This one has ten electrons. How many neutrons are present, in this oxygen? I'll let you try to answer that yourself. Look on your periodic table. Great. I see that all of you were able to figure out that oxygen has eight neutrons. At least the most abundant isotope of oxygen does. We'll talk more about isotopes in a minute. Finally, let's do another example; how about Yttrium 3+. Yttrium's a metal. I'll give you a hint. Have you been able to find it? OK. Here's another hint. It's atomic number is 39, so it has 39 protons. Now, this particular atrium has a charge of three plus. So how many electrons must this atrium contain? Wonderful, I'm glad you were able to figure out that this atrium has only 36 electrons, hm. Finally, how many neutrons does it contain? Well, in this case, you have to look out for, remember, you have to look up the atomic mass. Yttrium's atomic mass is 89. So, if the mass is 89, and I subtract the 39 protons, that would give me a total of 50 neutrons. There they are. I hope doing these examples has helped you figure out how to determine how many electrons and neutrons there are in a given species. Let's do a little bit of review before we dive into isotopes. You probably remember from one of the previous lectures, that all atomic masses are based on carbon-12 as a standard. We, as humans, arbitrarily decided to do that. We decided that carbon that has six neutrons and six protons has a mass of exactly 12. Everything else is relative to the carbon. But if you have a pretty decent periodic table and you look at it, you might notice, that the average atomic mass reported even for carbon which is our standard, is not exactly 12.00000. In fact it's 12.01 atomic mass units for the average atomic mass. Remember, that's the average atomic mass of all of the carbon on Earth. So where did the extra 0.01 atomic mass unit come from? Remember, if I had a mole of carbon atoms, that would be an extra 0.01 grams. Where did that come from? Well, that comes from the fact that not all carbon atoms have six neutrons. Some carbon atoms have seven neutrons and other carbon atoms have eight neutrons. And the property of different atoms having different types of neutrons is called isotopes. Isotopes are very useful for humans in many ways. Let's start by doing some calculations with isotopes, and lets start with carbon since we were just giving carbon as an example. Isotopes are atoms of the same element. So, they have the same number of protons, but there are different in their number of neutrons. So, for example, I might have two different isotopes of carbon. There are indeed two stable isotopes of carbon that are commonly found on Earth. One of the carbons is called carbon-12. It has a mass number of exactly twelve. So, it's got six neutrons and six protons. But there's another stable form of carbon called carbon-13. And you might have heard of yet another type of carbon called carbon-14. Carbon-14 gets used in radioactive dating of archaeological samples but that one's not actually stable, so since we're saying stable isotopes, I'm going to cross Carbon-14 out. Let's do our calculation using just the isotopes that don't decay over time, Carbon-12 and Carbon-13. Well, here are the symbols of Carbon-12 and Carbon-13. Both of them have 6 protons. So, that's shown here. The Carbon-12 has a mass of exactly 12, so this is the mass number. And the Carbon-13 has a mass of exactly 13, this is the mass number, you might recall from earlier. To calculate the number of neutrons, recall that we just take the mass number and subtract the number of protons, because the electrons don't have, a mass that matters. The mass of the electrons is negligible. So, Carbon-12 then, has six neutrons, and Carbon-13 has seven neutrons. So, there's two different isotopes of carbon, that are stable on Earth. Most of the carbon on earth has six neutrons. But there's a little bit of carbon on earth, about 1% of the carbon on earth, and that is in the plants we eats, in the carbon dioxide in the air, has this extra neutron and weighs 13 atomic mass units instead of 12. We can calculate the average atomic mass this way. We take the relative abundance of carbon 12. So lets suppose we had 100 carbon atoms. If 99% of the carbon atoms are carbon-12 then the relative abundance of carbon-12 is 99 divided by 100 which is 0.99. So, this is the relative abundance. We multiply that by the mass of carbon-12, which is exactly 12. Out of my 100 carbon atoms, I also have one Carbon-13. So, the relative abundance of Carbon-13 is 1 divided by 100 which is .01. I multiply that by 13 and then I add those two numbers together and I get the average atomic mass, which is 12.01. Pretty easy calculation to do. The average atomic mass is what's reported on the periodic table. Average atomic masses do vary a little bit from place to place on the face of the planet Earth. And that can be very useful sometimes when we're trying to determine the origin of a sample. In fact, the relative amount of Carbon-13 in a sample depends not only on where it was found on Earth but what time of day it is, if it is a plant. When plants are actively breathing in carbon dioxide during the day they have a different amount of Carbon-13 in their system than they do at night when they're resting. Isn't that interesting? Carbon-14, for example, is not stable, but decays over time. It has a type of radioactive decay. We can use the amount of Carbon-14 in the sample to determine how old the sample is because once a species is no longer animate, it's no longer consuming any other Carbon-14, samples like eating food, and so the amount of carbon 14 in that sample decays extremely slowly over time. The chemist can use that to determine how long the sample has been inanimate. So, this is used all the time to date things like mummies or plant fossils that are found. Another time that isotopes comes in handy, this is part of the why do I care portion of chemistry class Is when we're trying to determine if someone has taken some artificial hormones. For example, one of the things that can make you build muscle mass more quickly is having testosterone in your system. Testosterone is a naturally occurring steroid, that your body can synthesize and men have more testosterone in their system, generally, than women do. As it turns out, you can look at a sample of someone's blood and determine if the testosterone in the sample is testosterone that their body manufactured because then it has a certain ratio of Carbon-12 to Carbon-13. Or, is it testosterone that was manufactured in a laboratory, not by a human body? That sample has a slightly different ratio of Carbon-12 to Carbon-13. Now, we had a student at Duke many years ago, Derrick Lowe, who wrote a wonderful blog post about this. Around the time that Floyd Landis got in trouble for taking artificial testosterone, when he had the fastest time in the Tour de France, which happened in 2006. Now, shortly thereafter they stripped him of the title. They determined that some of the testosterone in his bloodstream was not testosterone that his body had created and they did that using mass spectrometry to look at the ratio of Carbon-12 to Carbon-13. Isn't that interesting? If you are really interested in that, I encourage you to read more. You can read more on, on Dr. Lowe's blog or you can read more on other sites, but it is really interesting. Let's do another example of isotopes. We've talked a lot about carbon, but carbons, most of the carbon on earth is one isotope. Most of it's Carbon-12. 99% is Carbon-12. There are other types of elements that have a bigger split between different types of isotopes. For example, chlorine. Chlorine exists in nature as Chlorine-35 and Chlorine-37. Chlorine-36 isn't very abundant. Isn't that strange? We don't really know why that's true, but we know that if we go out into nature and we take samples of things that have chlorine some of the Chlorine has a mass of 35 atomic mass units and some of the chlorine has a mass of 37 atomic mass units. Remember, this is the mass number of a specific atom. So, if we wanted to determine the number of neutrons. Well, we know the number of protons because that's the atomic number. We can just look that up on the periodic table. For chlorine it's 17. If we take 17 and subtract it from 35 we can determine that Chlorine-35 has 18 neutrons. Chlorine-37 then, has two more neutrons. It's got 20 neutrons. It's pretty simple math to do. In nature, the lighter isotope of chlorine is again more abundant, but this time the abundance isn't 99%, it's 70%. So, these are things we've determined by making observations of the world around us. The abundance of Chlorine-37 is 30%. Well, that makes the calculation for the average atomic mass of chlorine a little strange. It's the same calculation that we did before. I'm going to take the average abundance. If I had a 100 and 70% of them are Chlorine-35 then the abundance of chlorine-35 is 0.7. That's 70 divided by 100, right? And this is the mass Chlorine-35. I have 30% Chlorine-37, so I take 30 divided by 100. That gives me 0.3. Multiply that times the mass of Chlorine-37. I add those two numbers together, and I get that the average atomic mass of chlorine is 35.45, which I've rounded here to 35.5. And of course if I had a mole of chlorine, I would have 35.5 grams for that mole. Because I would have some Chlorine-35 and some Chlorine-37 in my sample. So, that's what you'd see in the periodic table, the average atomic mass, based on the relative abundance, abundances of the stable, naturally occurring isotopes. This has really been a great time, learning about isotopes and ions. Let's review what we've learned. For example, we've learned that the identity of the element is determined by the number of protons. That's something from a previous lecture. We've learned that neutral atoms have equal numbers of protons and electrons. So, if something is neutral we should have equal numbers of protons and electrons. In this lecture we learned about ions, in which case the number of protons and electrons is not equal. Anions have more electrons than protons and cations, which are positively charged, have less electrons than protons. So, the charge on the anion is negative, and the charge on a cation is positive. Again, for review, we know that isotopes have the same number of protons, but they differ in their number of neutrons. And we can calculate the average atomic mass of a sample of a certain type of atom if we know the relative isotopic abundances. We can write an equation for that.
