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Hey everyone, Dave Farina here. To get ready for the AP chemistry exam, we’re going to have to understand a lot of different terminology and concepts, and be able to answer fairly complicated questions about them. Because there are so many terms and concepts to know, let’s go through them all together in as condensed a format as possible, with one comprehensive video for each of the nine units the exam will cover. Unit 1 is on atomic structure and properties, so
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these will be all the very basics. We will discuss moles and molar mass, mass spectroscopy, pure substances and mixtures, electron configurations, photoelectron spectroscopy, periodic trends, valence electrons, and ionic compounds. We will be focusing on the more complicated topics as they arise but it will be a good idea to quickly review the fundamentals, so let’s start with an abbreviated version of some basic concepts.
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An atom is the smallest unit of an element, which are these things we see on the periodic table, such as copper. If we cut a copper rod in half, we get two smaller copper rods. We can continue doing this many times until we finally reach the tiniest thing that could be considered copper, and that’s a copper atom. Every element is made of a different type of atom which means it has a different number of protons inside, and if we were to break apart this copper atom,
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we wouldn’t have copper anymore. Atoms join together via chemical bonds to form molecules. Sometimes molecules can be considered elements, if they are comprised of only one type of atom, like diatomic oxygen and nitrogen. But when atoms of more than one element come together to form a molecule, we would call this a compound. One carbon atom and two oxygen atoms makes carbon dioxide. One oxygen atom and two hydrogen atoms makes a water molecule.
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A pure substance can be an element or compound. If we have a block of sodium metal, which is an element, that’s a pure substance, because it is made of just one thing, sodium atoms. If we have a sample of water, which is a compound, that’s a pure substance, because it is also made of just one thing, water molecules. Each water molecule has more than one element in it, hydrogen and oxygen, but it is still a water molecule, and the properties of water molecules
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are what determine the macroscopic properties of water. When we put more than one element or compound together, we get a mixture. Notice that we have two different substances in this mixture, with totally separate identities and properties. So a compound is a combination of atoms of different elements to form a single molecule, where all the atoms are chemically bound, like with water. But with a mixture, we have substances that are totally different elements or compounds,
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simply mixed together in the same space. Mixtures can be either homogeneous or heterogeneous. A homogeneous mixture involves substances that, when mixed together, are distributed evenly in the container. For example, sucrose dissolved in water. When you put sugar in water and stir, it will disperse evenly within the water, such that if you were to zoom in on any portion of this mixture, it would look exactly the same, with sugar and water molecules moving around
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one another. By contrast, a heterogeneous mixture will not have its components spread out evenly, so every section will not look the same. An example of this is a mixture of oil and water. We can clearly see that they will form distinct layers, such that if we zoom in on one part, we see only oil. If we zoom in on another part, we see only water. And if we zoom in at the interface, we see oil and water molecules up against each other but not mixing.
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So to summarize, we have pure substances and mixtures. A pure substance can be either an element or a compound. A mixture is made when more than one type of pure substance is mixed together, and this can be either homogeneous or heterogeneous. Next we are going to want to review atomic structure, that is the arrangement of protons, neutrons, and electrons to form atoms, and the parameters they determine.
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Looking inside an atom, the protons and neutrons sit in the nucleus, with the electrons existing far away from the nucleus. The proton and neutron have roughly the same mass, about 1.67 x 10-24 grams, which is a trillionth of a trillionth of a gram, while the electron has nearly 2,000 times less mass than that, at around 9.1 x 10-28 grams. And in terms of electric charge, the proton and electron each hold the fundamental unit of charge, which is 1.602 x 10-19 coulombs,
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though for the proton that will be positive, and for the electron it will be negative. The neutron is neutral meaning it has no charge. Each element is defined by the number of protons in its nucleus. For example, a carbon atom has six protons in its nucleus. Every carbon atom in the universe has six protons in its nucleus, and every atom that does not have six protons is not carbon. This is fundamentally how we define carbon. So carbon has an atomic number of six,
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because the atomic number of an atom is equal to the number of protons in its nucleus. Every element has its own unique atomic number, and therefore characteristic number of protons. The mass number of an atom is equal to the sum of the numbers of protons and neutrons, since each nucleon has a mass of approximately one atomic mass unit. However, although every element has a particular atomic number, elements do not have a specific mass number,
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because unlike the number of protons, the number of neutrons in the nucleus can vary for any given element. For example, carbon typically has six neutrons, which combined with the six protons will produce a mass number of 12. But it can also have seven neutrons, for a mass number of 13, or eight neutrons, for a mass number of 14. Atoms of a given element with different numbers of neutrons are called isotopes of a given element, so carbon atoms will always have six protons,
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but different isotopes of carbon can have six, seven, or eight neutrons, which correspond to different mass numbers. The atomic mass of an element, which is the other number on each block of the periodic table, is the average of all the naturally-occurring isotopes of that element with respect to their relative abundance. So it’s essentially the average mass of an atom of that element. Most carbon is carbon-12, with a little bit of carbon-13, and trace amounts of carbon-14,
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therefore the average mass of all carbon atoms will be just a tiny bit above 12, as this equation determines. We just multiply each mass number by a fraction of one representing that isotope’s relative abundance, and add them all together. This can be done for any element. Finally, since protons are positively charged and electrons are negatively charged, a neutral atom will have the same number of protons and electrons. If an atom gains electrons, it will
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become negatively charged, because the number of negatively charged particles will outnumber the positively charged ones. And if an atom loses electrons, it will become positively charged, because the number of positively charged particles will outnumber the negatively charged ones. When an atom is not electrically neutral, it is called an ion, either a cation or anion if positive or negative, respectively. With both isotopes and ions understood, we can
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understand how the masses of existing isotopes and their relative abundances can be determined for any element using a technique called mass spectroscopy. This instrument takes a sample, vaporizes it, ionizes it, and sends it through a tube where it is subjected to an external magnetic field. The particles then have their paths deflected by a degree that depends on their mass to charge ratio, and this information is received when they collide with a detector. A mass spectrum
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displays this mass to charge ratio, which for ions with a singular charge is essentially just atomic mass, against their relative abundance. So with a spectrum like this we can clearly see the naturally-occurring isotopes for this particular element with their respective mass numbers on the horizontal axis, and their relative abundances on the vertical axis. This is how this information was initially determined for every element. To quickly summarize, the atomic number of an atom
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is equal to the number of protons in the nucleus. The mass number of an atom is equal to the number of protons plus the number of neutrons in the nucleus. This means that we can calculate the number of neutrons in any atom by finding the mass number minus the atomic number. And the electrical charge on an atom is the number of protons minus the number of electrons. We can report these values using something called a nuclide symbol. These consist of the chemical symbol for the
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element, with the atomic number in subscript to the left, the mass number in superscript to the left, and the charge in superscript to the right. We will have to interpret these nuclide symbols, so just remember the definitions of these terms and where they go. For example, look at this nuclide symbol for magnesium. We see 12 down here, which means 12 protons, which is actually slightly redundant, because every magnesium atom in the universe must have 12 protons to qualify
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as a magnesium atom. Then we see 24 up here, which means there must be 12 neutrons, since 12 plus 12 is 24. And then this plus two charge means the atom has lost two electrons compared to the neutral atom, which leaves 10 electrons. With the submicroscopic structure of the atom understood, we need to be able to talk about matter on the macroscopic level, as well. This is hard, since molecules are way too small to see, and macroscopic amounts of a substance contain
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an unbelievable number of molecules. That’s why we came up with the concept of the mole. A mole is just a word that describes a number, like the way a dozen means 12. But it is a very large number that allows us to convert between atomic mass and grams. For example, carbon weighs on average about 12 atomic mass units. According to the definition of a mole, a mole of carbon atoms will therefore weigh 12 grams. This demonstrates how the mole is our
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way of converting between atomic mass units and grams, so that we can discuss molecules in terms of numbers, but have that number be so large that it represents a quantity that we can see with our eyes and do chemistry with. In other words, this way we can weigh out matter in grams and do chemistry with it, but still be talking about numbers of molecules and thus respect the ratios in which these molecules react. The number of items in a mole is called Avogadro’s number,
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which is equal to 6.022 x 1023, which is nearly a trillion trillion. So that’s precisely the number of carbon atoms in 12 grams of carbon, because one carbon atom weighs 12 atomic mass units. The mass of one mole of a substance is called the molar mass. Elements will have a molar mass equal to their atomic mass but in grams per mole instead of atomic mass units. Compounds will also have a molar mass, and it will be equal to their molecular mass,
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so to find the molar mass of a compound we simply add up the atomic masses of all the atoms in the molecule, and then we express that number in grams per mole instead of atomic mass units. It should be very easy to convert between grams and moles for any compound. Let’s say we want to know the number of moles in 28.35 grams of glycine. We can simply find the molecular mass by adding up the atomic masses of all the atoms in the molecule. Using the molecular formula of
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C2H5O2N, we get two times twelve for the two carbons, five times one for the hydrogens, two times sixteen for the two oxygens, and fourteen for the lone nitrogen. Adding those values up will give us a mass of 75, which if expressed in atomic mass units will be the molecular mass, but if expressed in grams per mole it will represent the molar mass, or the mass contained in one mole of glycine molecules. Then we convert our mass into moles.
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We will multiply our gram value by this conversion factor, putting moles on the top and grams on the bottom so that grams cancel, and doing the arithmetic we will get 0.38 moles of glycine. We can go the other way as well, from moles to grams. Let’s say we are looking at vitamin C, which has the molecular formula C6H8O6. Say we have 1.42 x 10-4 moles, but we need this in grams. Once again, adding up the atomic masses of all the atoms in the molecule,
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we can get a molar mass of 176 grams per mole. If we multiply our value in moles by this conversion factor, we can see that moles cancel, and we will get an answer of 0.025 grams of vitamin C. With moles understood, we can start to better understand compounds and their compositions, as well as how these compositions can be determined. This will involve discussing empirical and molecular formulas. The percent composition of a compound is the
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percent of the molecular mass that is represented by each element in a compound. This is easy to calculate if we know the molecular formula of a compound, because we can then know the molecular mass. If we know the molecular mass, we can just find the fraction of the molecular mass that is contributed by each element. Let’s say we want to know the percent composition of ammonia. We know that ammonia has a molecular mass of 17 atomic mass units, because the nitrogen atom
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has a mass of 14 atomic mass units, and each hydrogen atom has a mass of 1 atomic mass unit, for a total of 17. We can simply calculate the mass of each element present in the compound over the total mass of the compound to get the percent composition of the compound. If one nitrogen atom is 14, then 14 over 17 will give us 0.82, which times 100 gives us 82%. This means that the nitrogen atom in ammonia represents 82% of the mass of the molecule. The three hydrogen
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atoms have a total mass of 3, and 3 over 17 gives us 0.18. Multiplying by 100, that gives us 18%, so hydrogen represents 18% of the mass of the molecule. And 82% plus 18% does add up to 100%, so these calculations do make sense. This line of thinking is actually a great way to determine the molecular formula of an unknown compound. We can do this by first figuring out the empirical formula of a compound, which is the lowest whole number ratio of the number
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of atoms of different elements in a compound. Let’s say we combusted an unknown hydrocarbon, which is a compound consisting of only carbon and hydrogen, and collected the resulting carbon dioxide and water. After performing some basic calculations, we determine that there was 1.71 grams of carbon and 0.287 grams of hydrogen in the initial sample. Since these values are in grams, they do not tell us anything about the empirical formula, because every element has a different
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mass. Instead, we must convert these to moles to make sense of a numerical ratio. We can use the molar masses of each element to convert to moles. 1.71 grams of carbon times 1 mole over 12.01 grams gives us 0.142 moles of carbon atoms in the original sample. Doing the same thing for hydrogen, 0.287 grams of hydrogen times 1 mole over 1.008 grams gives us 0.284 moles of hydrogen atoms in the original sample. Let’s divide both of these numbers by the smaller number so that
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we can try to get a whole number ratio. 0.142 over 0.142 gives us 1, and 0.284 over 0.142 gives us 2, or a 1 to 2 ratio. So we can see from these calculations that there must have been twice as many hydrogen atoms as carbon atoms in the sample. This makes the empirical formula for the unknown substance, CH2. We must realize that this is not the molecular formula, which tells us the actual number of atoms of each element in the compound. The compound could have
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many more than one carbon atom, but however many carbon atoms are in the compound, there must be twice as many hydrogen atoms. At any rate, we can perform a calculation like this for any compound containing any combination of elements, we just use the molar mass of each element and convert the mass into moles to find the molar ratios, and therefore the empirical formula. We can also get the molecular formula if we have the molecular mass of the compound, which we can
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get through mass spectrometry. In such a case, we would just find out how many multiples of the formula unit are required to get a total mass equivalent to the molecular mass. Given the previous example with an empirical formula of CH2, let’s say that we knew the molecular mass was 42. The mass of CH2 is 14, 12 from carbon and two from hydrogen, and 42 divided by 14 is three, so we just multiply the formula unit by three to get C3H6, and that must be the molecular formula, as
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it obeys the ratio of the empirical formula, and has a mass that is equal to the molecular mass.
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With empirical and molecular formulas understood, we have to dive back into the atom and learn more about electrons. Precisely how are these distributed within an atom? This will be important to understand in order to discuss chemical reactions. Again, atoms contain both protons and electrons, and these have opposite charge, which means they are attracted to one another. This attraction is described mathematically by Coulomb’s law,
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which says that the force between two charged particles is proportional to the product of their charges divided by the square of the distance between them. If opposite charges this will be an attraction, if the same charge it will be a repulsion. Greater magnitude of charge means greater force, and closer together means greater force, while farther away means lesser force. Electrons themselves reside in things called atomic orbitals, or three-dimensional regions
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of probability surrounding the nucleus where an electron can be found, and there are different kinds of quantum numbers which will describe these orbitals. The first quantum number is the principal quantum number n. This refers to the energy level or the shell that the electron resides in. A higher n value means a higher energy and further away from the nucleus. The next number will be the angular momentum quantum number, L. This can have any value from 0 to n-1,
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meaning if n is 1, L is 0. If n is 2, L can be 0 or 1, and so forth. L will define the type of orbital the electron is in. An L value of 0 corresponds to s orbitals. Those are spherical, and they increase in radius as n increases. If L is 1, we are discussing p orbitals, which are lobes that extend on each of the X, Y, and Z axes. If L is 2, we are looking at d orbitals, which look a bit stranger. S, p and d are the most important ones for our purposes. Next,
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we have the magnetic quantum number, m sub L. This can be anywhere from –L to L, so if L is 2, and we are discussing d orbitals, L can be -2, -1, 0, 1, or 2. This is why there are five d orbitals per energy level, because there are five possibilities for m sub L and each one corresponds to an individual orbital. For precisely the same reason, there are three p orbitals per energy level, with m sub L values of -1, 0, and 1, and there is only one s orbital per energy level,
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with an m sub L value of zero. Lastly, there is the spin quantum number, m sub s. This will be positive one half or negative one half, and since a maximum of two electrons can fit in any atomic orbital, each pair will receive opposite spin values, which we can call spin up or spin down. The key thing to understand is that the n value describes a shell of electrons, and the L value describes a subshell. So there is an n = 3 shell, and within that there is a 3s
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subshell, and a 3p subshell, and a 3d subshell. Then m sub L describes an individual orbital within a subshell, and m sub s differentiates between the two electrons within an orbital. Now we need to understand how electrons fill up these orbitals. As n increases, the energy of the orbital increases, as we are moving farther away from the nucleus, so Coulomb’s law says the attraction to the nucleus will decrease. We should also know that within a shell
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the energy increases from s to p to d orbitals. So 1s is the lowest energy orbital, then 2s, 2p, 3s, 3p, and so forth. But this pattern isn’t followed precisely when we get to larger atoms, the first deviation being that the 3d orbitals are higher in energy than the 4s. Looking at this diagram, we can see the precise order of the orbitals in terms of increasing energy. Since a system will always want to be at the lowest energy possible, this is the order, from bottom to top,
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that an atom will arrange its electrons. This order in which the orbitals are filled is called the Aufbau principle. Additionally there is Hund’s rule, which says that when looking at a set of degenerate orbitals, which means orbitals of the same energy, as a set of p orbitals or a set of d orbitals will always be, we must place one electron in each orbital first before doubling them up. So for these p orbitals, each one gets a spin up electron first, and then we start
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generating pairs by placing spin down electrons. Each electron within an atom must be assigned an orbital, and the specific arrangement of electrons amongst the orbitals within an atom is called the electron configuration of the atom. Many properties of an element will depend on its electron configuration, so let’s make sure we understand these as well. The convention for reporting an electron configuration is to list all the types of orbitals
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that are occupied along with a number to indicate the occupancy of those particular orbitals. Each item in an electron configuration should have the n value, followed by the letter that corresponds to the type of orbital, and a superscript that describes the number of electrons contained in that subshell. So this would be read 2p4, which refers to a total of 4 electrons that exist in the 2p orbitals. Let’s make sure we understand that neither of these two numbers is telling us how
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many orbitals are being described, as that number is implied, since each energy level contains 1 s orbital, 3 p orbitals, and 5 d orbitals. So when assigning an electron configuration we are starting with the lowest energy orbital, the 1s, and building up to the higher energy orbitals according to the Aufbau principle and Hund’s rule, until all the electrons are assigned to an orbital. A convenient way to follow the Aufbau principle is to simply know what sections
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on the periodic table correspond to which subshells. Looking at the periodic table now, we must understand that each period, or row on the table, represents a shell, or a particular n value. The first row is n = 1, then n = 2, and so forth. Then, we must know that this section containing groups 1 and 2 is called the s block. This section is the p block. The transition metals are the d block, and the lanthanides and actinides are the f block,
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though we won’t be too concerned with f orbitals here. The only trick is that the d block is always one behind the period number in terms of principal quantum number. For example we can see that in the 4th period it’s actually the 3d orbitals that follow the 4s. If we can internalize this way of looking at the periodic table, then the Aufbau principle reveals itself as we simply read left to right and up to down on the table. Starting at the top left corner, the order would be 1s,
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2s, 2p, 3s, 3p, 4s, 3d, 4p, and so forth, which is precisely the order dictated by the Aufbau principle. This will make it much easier for us to assign electron configurations to any atom. We should now be able to assign the electron configuration of any element. This is easy to do if we simply look at where an element sits on the periodic table, and list off the subshells that the element will utilize by going left to right and up to down on the table until we get
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to the element in question. Take chlorine for example. In order to get to chlorine, let’s start up at the top left corner, and read off the subshells. 1s will be full, so 1s2. Same with 2s, so 2s2 and 2p6 from period 2. Then 3s2, and when we get to the 3p orbitals, we count 1, 2, 3, 4, 5 to get to chlorine. 3p5. That gives us 1s22s22p63s23p5 as the electron configuration for chlorine. This makes perfect sense, as neutral chlorine has 17 electrons to place,
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and this is the lowest energy configuration for the distribution of these 17 electrons. With electron configuration understood, we can begin to truly understand the structure of the periodic table, and why elements are arranged in the groups as they are. Let’s take a closer look at the periodic table and see what else we can learn from it. Looking at the table now, we can see rows called periods and columns called groups, and elements
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will be in the same group because they have similar electron configurations. Specifically, they have the same number of valence electrons, which are the electrons in the outermost shell. In group 1, all the configurations end in s1. In group 2, they end in s2. You can see that in every group, the electron configurations end the same way as the other elements in that group. So as we move forward and learn about the periodic table, it is the number of valence electrons that will
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determine the reactivity and properties of any particular element, as the valence electrons are the ones that are available to do chemistry. Those electrons that are not in the outermost shell, and are therefore not valence electrons, are called core electrons. These are the ones in the inner shells, which do not participate in chemistry. So most elements have many core electrons, and just a few valence electrons, particularly as we get lower on the table.
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We must comprehend a set of periodic trends, meaning properties that atoms possess which change in a predictably periodic way as we move in some direction along the table. The first property we will look at is the atomic radius. It’s difficult to measure the radius of a lone atom, so the convention is to tabulate lists of covalent radii, which are defined as one half the distance between
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the nuclei of two identical atoms that are bonded to each other. This should be roughly the same as the radius of the atom outside of the context of a chemical bond as well. When we examine various covalent radii, we notice that the radius will increase as we go down the periodic table. This is because when we go down a row on the table, we are increasing the n value by 1, thus adding a shell and placing the valence electrons farther away from the nucleus. This
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makes the atom larger, as well as its covalent radius. Then moving horizontally, as we move to the right along a period, the covalent radius will decrease slightly. This may seem counterintuitive, as the addition of electrons doesn’t seem like it should result in a smaller radius, but we must realize that as we move to the right we are also adding protons, given that the atomic number is increasing, and the more protons there are in the nucleus, the greater the electromagnetic
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attraction that will pull the electrons in the existing shells a bit closer to the nucleus. There are a few deviations to this trend, particularly if we were to examine the transition elements of a particular period, but in general, atomic radius will decrease as we move through a period. So radius increases going down and left, and decreases going up and right on the table. So that covers covalent radii, which essentially refer to the size of an atom. But what happens
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to this radius when an element loses or gains electrons to become an ion? Any change in the number of electrons should affect the radius in some way. As it happens, any time an atom loses an electron, the remaining valence electrons will still feel the same effective nuclear charge, but distributed amongst fewer electrons, so it will cause the radius to contract. This means that any cation has a smaller ionic radius than the covalent radius of the neutral
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atom. This difference can be dramatic if all the valence electrons are lost, since this will result in the removal of an entire shell, dropping down to the shell below. But if an atom gains electrons, the effective nuclear charge will be distributed amongst more electrons than before, and there will be additional electron repulsion amongst the valence electrons, which results in an expansion of the radius. We can see a concrete example with aluminum. If all
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three valence electrons are lost, it will lose its entire valence shell, and therefore will have a dramatically reduced radius. Sulfur on the other hand, when it becomes the sulfide ion, the two additional electrons will cause the radius to increase quite a bit, since there are no additional protons to pull the electrons, just more electrons that will push the radius out. We might sometimes compare isoelectronic species. These are atoms and ions that have
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the same electron configuration. For example, let’s look at the different species that can exhibit the electron configuration 1s22s22p6. We can see that beyond just neon, a number of different ions can have this configuration if the corresponding neutral atoms gain or lose a certain number of electrons. When comparing isoelectronic species, the radius will decrease as the atomic number increases. This is because they will all experience the same amount of electron repulsion,
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since they have the same number of electrons, but as we add more protons to the nucleus, this pulls the electrons closer to the nucleus. Next let’s look at ionization energy. This is defined as the energy required to remove the outermost electron from an atom in the gas phase and in its ground state configuration. The higher the ionization energy, the more difficult it is to remove the electron, which tells us something about the atomic radius of the atom as well as
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the effective nuclear charge felt by the electron. Each element will have a first ionization energy, which is the energy required to generate the 1+ cation, and they will also have successive ionization energies, like the second ionization energy, to go from 1+ to 2+, and so forth. Each ionization energy will be larger than the last, because it will get harder and harder to remove electrons the more positive the ion becomes, as each ionization is a further destabilization.
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The electron that is removed will always be the outermost electron. As the atom gets larger, the outermost electron gets farther away from the nucleus, and therefore becomes easier to remove. Every time we add a shell, we are moving further away from the nucleus, so ionization energy decreases as we move down the periodic table. Since atomic radius also decreases going to the right within a period, we can expect the ionization energy to increase at the same
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time. As we go, we are adding protons, contracting the radius, and holding electrons more tightly, so they are harder to remove. That means in general, while atomic radius increases down and left, ionization energy will increase up and right, precisely the opposite of the radius trend. That means helium is the most difficult element to ionize, with a single shell that is totally full and close to the nucleus, while francium is the easiest, with a lone electron in an outermost
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shell that is very far from the nucleus. The energies of the electrons in an atom can be determined by photoelectron spectroscopy. With this type of spectrum, the energy required to remove an electron from a particular subshell is shown on the horizontal axis, and then the vertical axis tells us how many of those electrons are in that subshell. Further to the left means a greater binding energy which means electrons that are closer to the nucleus, starting
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with the 1s electrons. Then moving to the right they get lower in energy and farther from the nucleus. When we see these spectra we should be able to recognize which peak corresponds to whch subshell based on its position on the horizontal axis, and also state how many electrons are in each subshell based on the height of the peak. We also want to learn about electron affinity. This is precisely the reverse concept of ionization energy, it is the energy change
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involved with adding an electron to a neutral atom in the gas phase, thus forming a negatively charged ion. This process could absorb energy or release energy, depending on the element, and a negative electron affinity will mean that the process is actually favorable for a given element. Looking at this table, we can see that the trend is similar to the ionization energy trend, since the harder it is to remove an electron, or the higher the effective nuclear charge,
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the easier it is to add an electron, and thus a greater electron affinity. This is why elements like fluorine and chlorine have very large electron affinities, as gaining an electron will endow them with noble gas electron configuration, which is a very stable situation. So in general, electron affinity increases going up and right on the table, with some exceptions. Noble gases do not follow this trend, as with a full shell of electrons, it is typically
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not favorable to add another electron, so we discount them when considering this property. And finally, let’s examine electronegativity. Electronegativity is a measure of how well an atom can attract electron density towards itself, which is measured by looking at the way electrons are shared in chemical bonds. The more strongly it can attract electrons, the greater its electronegativity. Electronegativity will depend on atomic radius, since a smaller atom with
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a greater effective nuclear charge will attract electrons more strongly, so the electronegativity trend will be the same as the ionization energy trend, it will increase going up and right on the periodic table. Fluorine will have the greatest electronegativity, and francium will have the lowest. Again, we will exclude the noble gases from this trend, as with their full valence shells, they are not likely to share electrons, making electronegativity meaningless for those
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elements. There is a common point of confusion that we should make abundantly clear. We must make the distinction between electronegativity and electron affinity, because the latter involves an actual ionization and an associated energy change that is measurable. The former just describes a relative calculation of how well an atom attracts the electrons in a bond, which does not involve any transformation, and it is listed on an arbitrary relative scale from zero to four. And
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so to put it all together, atomic radius increases going down and left on the table, while ionization energy, electron affinity, and electronegativity all increase going up and right on the table. And that concludes a review of Unit 1. I’ll see you over in Unit 2 for more chemistry.
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