Understanding Periodic Trends in Chemistry

In this video, we're going to go over periodic trends such as atomic and ionic radius, electronegativity, ionization energy, electron affinity, and metallic character. So let's start with atomic radius. The size of atoms increases as you travel in this general direction of the periodic table. That is, as you go to the left, the size of atoms increases, and as you go down... the atoms get bigger. So let's take two examples, hydrogen and helium. On a periodic table, hydrogen is to the left of helium, so that means that hydrogen is bigger than helium. But now, why is that the case? Because if you look at the atomic mass for helium in the periodic table, it has an atomic mass of 4. And for hydrogen, the atomic mass is 1. So how is it that hydrogen is bigger when helium is more massive? To understand this, we need to jump into physics for a moment. You know that light charges repel, and opposite charges attract. So if you have a negative charge and a positive charge, these two will feel a force of attraction. Now if you increase the number of charges, let's say if you have two protons and one electron, the force will be two times as large because the charge was increased. The more protons you have, the greater the force of attraction between the protons and electrons. The second factor to consider is distance. If you have a proton and an electron that are very far apart, the force that they feel will be relatively weak. But let's say if you cut the distance in half. So let's say currently they're about 4 meters apart. But if you bring the proton and the electron closer, let's say if they're 2 meters apart, If you decrease the distance by a factor of 2, the force of attraction between these two particles will increase by a factor of 4. So the electrostatic traction is much larger if the charges are brought closer to each other. So if you increase the charge... of any of the protons or electrons, the force of attraction increases. So the more protons that a nucleus has, the greater the force of attraction between the nucleus and the electrons. And the second factor is distance. If you increase the distance, the force of attraction will decrease. And if you decrease the distance, the electrostatic force of attraction will increase. Now, let's think about what this means. Inner core electrons that are very close to the nucleus, they will feel a very strong force of attraction between itself and the nucleus. Whereas electrons that are away from the nucleus, like valence electrons, the force of attraction between those electrons and the nucleus will be relatively weak, since they're further away. Now, hydrogen has only one proton in its nucleus, so the charge of the nucleus is plus one, and it has one electron. Now, helium has two protons in its nucleus, and it has two electrons. So because the nuclear charge of helium is greater than that of hydrogen, the force of attraction between those electrons and the nucleus is much larger in helium. As a result, the size of the... electron cloud shrinks as those electrons are brought closer to the nucleus. And that's one reason helium is smaller than hydrogen even though it has a greater atomic mass. Because the charge of the nucleus increases, the size decreases. So as you travel from left to right across the periodic table, the atomic size will decrease as the number of protons increases. Now, let's say if we were to compare lithium versus hydrogen. Which atom is bigger? Now, lithium has three protons and hydrogen has one. yet lithium is bigger than hydrogen. Why is that? I mean, don't lithium have more protons? Now, the trend as you go from left to right, the size decreases because... because the number of protons increases, but the number of shells is the same. Here, hydrogen and lithium are not in the same row. They're in different rows. Here's hydrogen, here's helium, and lithium is below hydrogen. Now, for two elements in the same row, the one that's on the right is usually the smaller atom. But now, for two atoms that are in the same group, the one that's below it is usually the bigger one. Remember. atomic size increases as you go down. So now let's talk about why. Why is it that lithium is bigger than hydrogen? So hydrogen has one proton in its nucleus and it has one electron. Now lithium has an atomic number three, so it has three protons in its nucleus. Now in the first shell, it can only fit two electrons. These are core electrons. So, it has to put the third electron in the second shell. And because that electron is further away from the nucleus, the force of attraction between this electron and the nucleus is reduced. Remember, this is the force of attraction. If you increase the distance, the force of attraction decreases. So the reason why lithium is bigger than hydrogen is because it has two electron shells as opposed to one. And that's why as you go down the group in a periodic table, atomic size increases because for every row that you add, you add a new shell of electrons. And so the atom gets bigger. Now, there's another reason that... you need to consider and that is the concept of shielding. The inner two core electrons partially shield the outer valence electron from the nucleus. Now this electron is attracted to the three protons in the nucleus. As we said before whenever you have opposite charges they attract each other. So what if you have like charges let's say two electrons like charges repel each other. So, this electron is repelled by these two electrons. So, even though it feels a force of attraction between... itself in the nucleus, it's also repelled by the other electrons between it and the nucleus. So you have to consider also the shielding effects of the inner core electrons. Now let's work on some examples. Let's say if we have three elements, chlorine, magnesium, and phosphorus. Rank the following elements in order of increasing atomic size. So that's from small to large. So which of these atoms is the smallest and which one is the largest? So the first thing you want to do is you want to look at the periodic table. And you want to place them in order. Magnesium is in the third row of the periodic table. and it's in group 2. It's an alkaline earth metal. Phosphorus is also in the same role and it's in group 5a and chlorine is a halogen in group 7a. Now because these three are in the same row, this is going to be fairly easy. We know that the atomic size increases, or rather it decreases, from left to right. So that means that magnesium is bigger than phosphorus, and phosphorus is bigger than chlorine. So if we want to rank it in order of increase in atomic size, we need to put the smallest first, and then the largest last. So magnesium is bigger than phosphorus, which is bigger. Let's try another example. Let's say if we have calcium, beryllium, and strontium. Rank the following elements in order of decreasing atomic radii. So notice that these three elements are all alkaline earth metals. They're found in group 2A. So let's place them in order. So first is BE. Below BE is magnesium, and then below that is calcium, and then strontium. Now we know that atomic size increases as you travel down a group. Now we want to rank it in order of decrease in atomic size So we know that strontium is the largest Beryllium is the smallest and calcium is in between. So to put it in decreasing order we want to rank it from large to small. So we're going to start with strontium, then calcium, and then calcium is bigger than beryllium. And so that's how you can rank it in order of decrease in atomic size. Now let's try one more example. So this time let's say if we have four elements. Fe, Cs, Sulfur, and Helium. Rank the following elements in order of increasing atomic size. So the first thing you want to do is you want to place every element in their respective positions based on the periodic table. Cesium is an alkaline metal in group 1, but it's in the 6th row, so it's like over here. Fe is a transition metal, and it's in the 4th row, so it's somewhere in this vicinity. Sulfur is a chalcogen found in group 6a, and helium is a noble gas that's in the first row. So remember, atomic size increases as you go to the left and as you go down. So basically... it increases in this general direction, which tells us that cesium is the largest, then it's Fe, then it's S, and then it's helium. So we wish to rank it in order of increase in atomic size so that's from small to large so helium is smaller than sulfur and sulfur is less than Fe and iron metal is smaller than an atom of cesium so this is the answer now let's move on to ionic radii So consider an atom of lithium and a lithium ion. An atom is a particle that has an equal number of electrons and protons, so atoms are electrically neutral. Ions have unequal number of protons and electrons and so they have a charge. An ion with a positive charge means that there are more protons than electrons. So lithium is an atom with three protons, so the charge of the nucleus is three. And as we mentioned before, it has two electrons in its first shell, and one valence electron in its second shell. The lithium plus ion, which is also known as a cation, cations are positively charged ions. It has three protons, but it lost an electron. So basically, it lost an entire shell. So therefore, positively charged cations are significantly smaller than their parent atoms. Because they have less electrons and therefore less electron shells, they will be significantly smaller. Now what about anions, or negatively charged ions? How do the sizes of these ions compare with their parent atoms? So let's use nitrogen as an example, and the nitride ion. Nitrogen has seven protons in its nucleus. And an atom of nitrogen has seven electrons. So in the first shell, it's going to have two. And in the second shell, it's going to have 5. So nitrogen has 2 core electrons, or inner electrons, and it has 5 valence electrons. The valence electrons are the electrons in the outermost energy level. Now the nitride ion has 3 more electrons than protons. It has 7 protons and 10 electrons. Protons have a positive charge, electrons have a negative charge. If you add these two numbers, you're going to get a net charge of minus 3. So let's draw the picture for the nitride ion. Now the first shell contains two electrons, but the second shell has a total of 8 electrons, so that the total number of electrons is 10. So because of the extra three electrons, there's more electron repulsion, which causes the second shell to expand. So negatively charged ions, or anions, are significantly larger than apparent atoms. So make sure you remember this for ionic radii. Ions with a positive charge will be relatively small, and ions with a negative charge will be relatively large. Now the trend for ionic radii is very similar to that for atomic radii, meaning ionic radii increases as you go down and increases as you go towards the left. This is especially true for ions with similar charges, if you're comparing two ions with a positive charge or two ions with a negative charge. Let's look at the elements of ions in a row. The sodium ion is bigger than the magnesium ion. Both of these two ions are isoelectronic. They have the same number of electrons, but because magnesium has more protons than sodium, magnesium is going to be smaller. And aluminum, which has a plus 3 charge, is even smaller than magnesium. So as you can see, as the charges increases, the size of the ion decreases. And as you travel to the right, it decreases to the right, but ionic radii increases towards the left. Now as you jump from the positive ions towards the negative ions, the size will greatly increase. For example, phosphide, which has a negative 3 charge, it is significantly bigger than almost all of these positively charged ions on the left. And then the size decreases as you go towards the right. So the sulfide ion is smaller than the phosphide ion. And then the chloride ion, which, let me see if I can fit it here, is even smaller than the sulfide ion. So the size generally decreases as you go from left to right. But if you're comparing a negatively charged ion with a positively charged ion, generally speaking, the negatively charged ion will be bigger than the positively charged ion. Now what about as we go down a group? Ionic size increases as you go down. For example, the lithium ion is smaller compared to the sodium ion. The sodium ion is larger. And the potassium ion is even bigger than the sodium ion. So as you go down, the number of electron shells increases, and therefore the ions will increase in size. So if you compare the nitride ion with the phosphide ion, the nitride ion is pretty big, but the phosphide ion is even bigger. So ionic radii increases as you go down a group. Now let's try a problem. Let's say if we have the following ions. The fluoride ion, the magnesium cation, an atom of neon, the sodium cation, and the oxide ion. Now rank the following atoms and ions in order of increasing size. Now, to do that, we need to determine if any of these atoms or ions are isoelectronic, meaning that if they have the same number of electrons. If they're isoelectronic, then all you need to do to rank them in order of increase in size is look at the atomic number. For particles that have the same number of electrons, the ones that have more protons will be smaller in size. As you increase the amount of protons in the nucleus. The increased nuclear charge will cause the electrons to move towards the nucleus, decreasing the size of the atom. So let's count the electrons first. The number of electrons is equal to the atomic number minus the charge. So if we consider fluorine first, fluorine has an atomic number of 9 and a mass number of 19. So it's 9 minus negative 1. So fluorine has 10 electrons. Now what about magnesium? If you look at the periodic table, magnesium has an atomic number of 12 and a mass number of 24. The small of these two numbers is the atomic number. So for magnesium it's going to be 12 minus a charge of plus 2 which is 10. So magnesium has 10 electrons. Now for neon it's an atom. which has an atomic number of 10, a mass of 20, and for atoms they're neutral, so they don't have a charge. So it's 10 minus a charge of 0. Neon 2 has 10 electrons. Now what about sodium? Sodium has an atomic number of 11 and for oxygen has an atomic number of 8. So for sodium it's 11 minus plus 1 which is equal to 10 and for oxygen is 8 minus negative 2. oxygen has 10 electrons. So all of these particles are isoelectronic with each other. They have the same electron configuration. So now all we need to do is look at the atomic number. atomic number of fluoride is 9, magnesium is 12, neon is 10, sodium is 11, and oxygen is 8. And as you mentioned before, negatively charged ions are generally larger than positively charged ions. And one of the most protons, or the highest atomic number, is going to be the smallest one. So oxide. We want to rank it in order of increase in size, so let's start with the smallest one. Magnesium has the most or the greatest number of protons. So it's going to be the smallest. It's smaller than sodium and sodium is smaller than neon Which is smaller than fluoride and the oxide ion is going to be the biggest So let's draw the relative sizes. So this is the smallest. That's a little bit bigger. Then this is going to be bigger. And then the size just continues to increase. So as we can see, oxide has the least amount of oxygen. number of protons, so that's why it's bigger. Keep in mind, each of these atoms and ions, they have only two shells of electrons. So for particles that have the same number of shells, the one with the most protons will be the smallest. The one with the least number of protons, in this case oxide, is going to be the biggest. So let's compare the magnesium ion and the oxide ion. Just to get a better understanding of this concept. So we said that magnesium has an atomic number of 12, which means that the charge of the nucleus is 12, since it has 12 protons. And it has a total of 10 electrons. So here's the first shell, and here is the second shell. Which, I think I could draw a better circle. So the first shell contains two, and the second shell has eight. Now, oxide also contains ten electrons, but the charge is eight, since oxygen only has eight protons in its nucleus. So the oxide ion is going to be significantly bigger than the magnesium ion. And we can see why it's bigger. Because magnesium has more protons, the electrostatic force between the nucleus and the electrons is stronger because of the increased charge. And so magnesium, it shrinks in size. The nucleus, it pulls the electrons toward the nucleus, making the ion smaller. Now, in the case of oxygen, the... force of attraction between the outer electrons and the nucleus is relatively weaker because the charge of the nucleus is less. Plus, since the net charge is negative, there's a lot more electron repulsion and so the atom expands when you have extra electron repulsion. So to summarize what we just went over, remember negatively charged ions are usually bigger than positively charged ions. And for isoelectronic species, that's particles that have the same number of electrons. which is going to have the same number of electron shells, the ones with more protons will be smaller than the ones with less protons. So as you increase the nuclear charge, and if the number of shells remain constant, the particle size will decrease. That extra nuclear charge will cause the electrons to contract towards the nucleus. So now let's move on to electronegativity. Electronegativity is the ability of an atom to attract an electron to itself. Electronegativity increases towards fluorine. Fluorine is the most electronegative element in the periodic table. On the upper right corner of the periodic table you have the nonmetals and on the left side you have the metals. Metals tend to be electropositive. They like to give away electrons and form positively charged ions. Nonmetals like to acquire electrons and form negatively charged ions. So nonmetals tend to be electronegative. Electronegativity increases as you go up and as you travel towards the bottom of the table. towards the right. So fluorine is more electronegative than oxygen, and oxygen is more electronegative than nitrogen. So let's say if you have a fluorine atom, and if you add an electron to it, fluorine will turn into the fluoride ion. The nonmetal will become a negatively charged anion. The nonmetals tend to be electronegative. they like to acquire electrons. Now if you have an atom of sodium, sodium wants to give away an electron to form a positively charged cation. So metals tend to be electropositive. They like to give away electrons. So those are some things to know. Now let's say if you have four elements, silicon, magnesium, chlorine, and aluminum. Rank the following elements in order of increasing electronegativity. So, look at the periodic table and place them in order. So, each of these elements are in the same row. Magnesium comes first, it's in group 2. Then aluminum, which is in group 3A. And then silicon, that's in 4A. And chlorine is in group 7A. So electronegativity increases towards the right. So if we want to rank it in order of increasing electronegativity, that's from low to high. Magnesium is going to be the least electronegative, it's a metal. And then it's aluminum, which is also a metal. And then silicon, which is a metalloid. And then chlorine is the most electronegative, that's a nonmetal. So nonmetals are usually... So this is the answer. Now what about these? Let's say if we have tin, germanium, lead, and carbon. Rank the following elements in order of decreasing electronegativity. Now each of these elements are found in group 4A of the periodic table. So let's put them in order. First we have carbon, and then below that, GE. germanium and then tin metal and then lead. Electronegativity increases as you travel up a group so we want to rank it in decreasing order so we need to put the most electronegative element. We need to rank it from high to low. So the highest is carbon which is a nonmetal, then is germanium which is a metalloid, and then tin metal, and then lead metal. So the metals, they have the lowest electronegativity. So this is the answer. That's how you can rank it in order of decreasing electronegativity. Now, which element is more electronegative, silicon or nitrogen? If you place these elements in their respective position, silicon is in group 4A in the third row, nitrogen is in group 5A. in the second row. Electronegativity increases towards fluorine. So it increases as you go up and to the right. So therefore you can clearly see that nitrogen is more electronegative than silicon. Now what about nitrogen versus sulfur. If you place them in their respective positions, nitrogen is in group 5A in the second row, sulfur is in group 6A but in the third row. But electronegativity increases this way. So which one is more electronegative? So as you travel up, electronegativity increases, but as you travel to the right, it decreases. So sulfur is to to the right of nitrogen, but it's also below nitrogen. So how can you tell which one is more electronegative in this case? It turns out that the electronegativity increases more when you go up than when you go to the left. So the increase for going up one row is greater than a decrease from traveling one unit to the right. Now if you're ever unsure, you can always use the equation that I just showed you. So the increase for going up one row At this point, you can look at the election negativity table and see which one is higher. The election negativity for nitrogen, according to most textbooks, is roughly about 3.0. For sulfur, it's 2.5. So as you can see, going up has more priority than going to the right. So let's go over some common electronegativity values for elements like boron, carbon, nitrogen, oxygen, fluorine, phosphorus, sulfur, chlorine, bromine, and iodine. The election negativity value for boron is about 2.0. For carbon, it's 2.5. For anode, it's 3.0. For O, it's 3.5. And for fluorine, it's about 4.0. Now for phosphorus, it's roughly around 2.0. one. For sulfur 2.5 and for chlorine it's about 3.0. For bromine 2.8, iodine 2.5. So as you can see there's a large increase as you go from to rho 3. Here the increase is 0.9 and here it's about 1. So that's why nitrogen was significantly higher than that for sulfur. So if you need to go this way, generally speaking, this is the one that's going to win. But not always though. For example, bromine is higher than sulfur. So there are some exceptions. So you may just need to know some of these values if you ever get an unusual question like that. By the way, the electronegativity for hydrogen is about 2.1. So even though hydrogen is to the left of boron, boron is in row 2, hydrogen is in row 1. So hydrogen is a little bit more electronegative than boron. Now the next train that we need to talk about is metallic character metallic character increases in this general direction Towards the metals the nonmetals are located in the upper right corner of the periodic table and the metals is just to the left But metallic character increases as you travels to the left and down across the periodic table So let's go over some examples. So let's say if we have elements such as silicon sodium sulfur, aluminum, and chlorine. Rank the following elements in order of increase in metallic character. So as we've been doing before let's place it based on their respective positions in a periodic table. Now each of these elements are located in the same row in the periodic table. Sodium is found in group 1 and then aluminum in group 3a. silicon in group 4a, sulfur in group 6a, and then Cl in group 7a. So metallic character increases as you travel towards the left on the periodic table. So if we wish to rank it in order of, let's say, increasing metallic character, we need to rank it from low to high. So chlorine, which is a nonmetal, is going to have the lowest metallic character. Sulfur is also a nonmetal. Silicon is a metalloid, aluminum is a metal, and sodium is a metal. Now let's think about what this means. Sodium really wants to get rid of its electrons more than aluminum. Both of these elements are metals, but sodium has a greater metallic character than aluminum, so sodium is more electropositive. It wants to get rid of its electrons with a stronger force than aluminum. Now between sulfur and chlorine, chlorine is more electronegative, so it's less metallic than sulfur. Silicon is a metalloid. It's in between a metal and a nonmetal. So this is the answer. That's how you can rank it in order of increasing metallic character. Now metallic character increases as you go down a group. Let's consider the group 4A elements, like carbon. silicon, germanium, tin, and lead. As you go down the periodic table, the elements change from nonmetals to metals. Carbon is a nonmetal, silicon is a metalloid, and germanium is a metalloid. But tin is a metal and lead is a metal. So as you can see, metallic character increases as you travel down a group. Let's try one more example. Let's say if we have Gallium, Manganese, Nitrogen, Helium, and Francium. Rank the following elements in order of decreasing metallic character. Francium is located in group 1 in the 7th row. And then you have manganese, which is a transition element. And then gallium is to the right of that. And then there's nitrogen and helium. So metallic character increases in this general direction. So if we want to rank it and decrease in order, we need to rank it from high to low. So francium has the greatest metallic character. Then it's manganese, and then Ga, and then N. and then He. The next topic that we need to talk about is ionization energy. Ionization energy is the energy required to remove an electron from a gaseous atom. Now generally speaking, it's easier to remove an electron from a metal than it is to remove from a nonmetal. Metals like to give away electrons. So it doesn't require that much energy to remove an electron from it. Nonmetals like to acquire electrons, so it's harder to remove an electron from a nonmetal. The trend for ionization energy is in this direction. Ionization energy increases towards helium. It increases as you go up across the periodic table and towards the right. So the main reason why it increases as you travel from left to right... is due to the increase in nuclear charge. From left to right, the number of electron shells, or the principal quantum number, remains the same. However, the number of protons in the nucleus increases. And as you add more protons to the nucleus, those protons will have a tighter grip on the electrons. So it requires more electrons, I mean, it requires more energy to remove those electrons. So let's compare lithium and beryllium. Lithium has three protons in its nucleus, and beryllium has a nuclear charge of four. Now keep in mind, beryllium is smaller than lithium because beryllium has a higher nuclear charge. So here's a question for you. Is it easier to remove the electron from lithium or a valence electron from beryllium? It's going to be easier to remove the electron from lithium. The ionization energy for lithium is about 520, but for beryllium it's about 899. So it requires much more energy to remove an electron from beryllium than to remove it from lithium. For one thing... The force of attraction between the electron and the nucleus is weaker compared to the force of attraction between an electron and beryllium than the nucleus. The nuclear charge of beryllium is much greater than that of lithium, so the nucleus has a tighter grip on this electron. So that's one reason why it's harder to remove it. But also, beryllium is smaller than lithium, and so the distance between the electron and the nucleus is much less than the valence electron and the nucleus within lithium. So remember, if you decrease the distance, the force of attraction increases. So because beryllium is smaller, and plus the fact that it has a higher nuclear charge, means that beryllium... holds on to that valence electron with a tighter grip than lithium. And so that's why it requires so much more energy to remove that valence electron in beryllium. It has to do with the increased nuclear charge and the smaller size. Both of those factors increases the ionization energy of BE. Now what about between lithium and sodium? Which element has a higher ionization energy? Sodium is below lithium and ionization energy increases as you go up. The ionization energy for sodium is about 495, but for lithium it's 520, so it doesn't vary much. Now lithium, as we mentioned before, has a nuclear charge of 3, and sodium has 11 protons. So sodium is bigger than lithium. It has three electron shells. So why is it easier to remove a valence electron from sodium than it is to remove it from lithium, considering that sodium has a higher nuclear charge? Now granted, as the charge increases, we know that the electrostatic force of attraction increases. So an atom with a higher nuclear charge means that it has electrons that are harder to remove. However, you have to take into account distance. The distance between this electron and the nucleus is significantly larger than the distance between this electron and the nucleus. So if you increase the distance, the force greatly decreases. And so because this electron is very far away from the nucleus, it's relatively easy to remove that electron. And that's why sodium has a slightly lower ionization energy than lithium. It's because of the increased distance between that electron and the nucleus, even though it has a higher nuclear charge. Now we can see that distance plays a greater role in affecting the force of attraction than the charge. Because the electron is further away, the ionization energy is less, even though the nuclear charge is greater. And this is consistent with Coulomb's Law. His law describes the relationship between the electrostatic force between two charged particles. So let's say if you have a proton and if you have an electron. The force of attraction between these two particles is called the force of attraction. particles is proportional to the charge and inversely related to the distance between them which is R so if you double the value of Q the force will double in value however If you double the distance, because it's squared, 2 squared is 4, the force will reduce by a factor of 4. So the distance between the charges has a greater impact than... the magnitude of the charge based on the equation. So to summarize what we've learned, as we travel down a group, the ionization energy decreases because the distance between the protons and the valence electrons increases. And as we travel to the right, ionization energy increases due to the increase in the effective nuclear charge. As you go to the right, the principal quantum number remains the same. So the number of electron shells is roughly about the same. Now there are some exceptions that you need to be aware of. For example, between beryllium and boron. The ionization energy for boron is 800 and for beryllium the first ionization energy is 899. Now boron is to the right of beryllium in the periodic table so typically we should expect that boron should have a higher ionization energy but it doesn't. The question is why? Now the last electron in beryllium is the 2s2 electron. However, the last electron in boron is the 2p1 electron. 2p, on average, is farther away from the nucleus than 2s. If you write the electron configuration, it's 1s2, 2s2, and then 2p6. But if you draw the orbital diagram energy levels, it looks like this. So 2p is higher in energy, which means that it's further away from the nucleus. And as we mentioned before, if you increase the distance between the electrons and the nucleus, on average, the ionization energy tends to decrease. And so... that's why we see a slight drop as you go from be to boron the same is true from magnesium to aluminum as you go from the s2 level to the the first P level there's a temporary decrease even for calcium to gallium so from magnesium to aluminum it changes from 735 to 580 which is pretty significant From calcium to gallium from group 2a to group 3a It drops from 590 to 570 so I mean 579 so the change between these two is not that great, but These are exceptions that you want to be aware of if you're tested on it Now, there are some other exceptions that you need to be aware of as well. It's not over yet. And that's going from nitrogen to oxygen, or phosphorus to sulfur, or arsenic to selenium. The first ionization energy for nitrogen is oxygen. is 1402 and for oxygen is 1314 at least according to the textbook I'm using now why does it temporarily decrease from n to O the electron configuration for nitrogen ends in 2p3 and for oxygen it ends in 2p4 Now for nitrogen, it has three unpaired electrons. But for oxygen, it has a paired electron. This paired electron is the one that's being removed. And due to the electron repulsion between the two electrons in this orbital, that allows this electron, the one being removed, It allows it to be removed with ease. Because that electron is being repelled by the other electron in that orbital, it doesn't take much energy to remove that valence electron. And that's why we see this drop in ionization energy. It's due to the electron repulsion between these two paired electrons. Now let's try some examples. Rank the following elements in order of increase in ionization energy. That is first ionization energy. So let's say if we have the elements gallium, bromine, potassium, chromium, and arsenic. So once again, you want to look at the periodic table, and you want to place them in their respective positions. So each of these elements are in the first, not the first row, but the fourth row. The first one is potassium, and then after potassium it's chromium, which is a transition metal. And then you have gallium, and then arsenic, and then bromine, which is a nonmetal. The first ionization energy increases as you travel left to right on a periodic table. So let's say if we want to rank it in order of increasing ionization energy, we could simply write it like this. So this is the answer. Now what about these three? Phosphorus. arsenic, nitrogen, and antimony. Go ahead and rank these elements in order of increasing or rather decrease in first ionization energy. Now these elements are all found in group 5a of the periodic table. So if we place them in order, first we have nitrogen, below that is phosphorus, then arsenic, and then after that, antimony. Now, the first ionization energy increases as you go up a group within the periodic table. So, if we want to rank it and decrease in order from high to low, we need to start with the highest, which is nitrogen. And that's greater than phosphorus, which is larger than arsenic, and which has a higher first ionization energy than antimony. So, that's how you can rank it in order of decrease in first ionization energy. Now let's try another example. Let's say if we have fluorine, phosphorus, helium, francium, and vanadium. So let's put this in order. So francium is all the way at the bottom, towards the left. It's in the 7th row, 1st column. Vanadium is a transition metal which is in this region in the periodic table. And then we have phosphorus, fluorine, and helium is to the upper right of fluorine. So first ionization energy increases towards helium. So therefore, let's say if we want to rank it in order of increase in ionization energy from low to high, let's start with francium, which is less than vanadium, that's less than P, which is less than F, and that's less than He. So it helps if you can place the elements based on where they're located in a periodic table, and then use the trend to rank them in the appropriate order. Now sometimes you may have questions about the second ionization energy and the third and so forth. The second ionization energy is the energy required to remove the second electron. That second electron may be a valence electron or maybe a core electron. So let's consider aluminum. Aluminum is in group 3a and it has a total of 13 electrons. Its atomic number is 13 so as an atom A neutral atom has 13 protons and 13 electrons. Now let's draw the aluminum atom. So the charge of the nucleus is 13. In the first shell, it's going to have 2 electrons. In the second energy level, it's going to have 8 electrons. So right now we have a total of 10. And in the third energy level, it has three valence electrons. Now the first ionization energy of aluminum is the energy required to remove the first valence electron. Let's say this one. So it's going to take about 580 kilojoules of energy to remove one mole of electrons from aluminum. And that's really one electron per one aluminum atom, but it's 580 kilojoules per mole. Now, for the second electron, the first ionization energy is higher. It's 1815. And it makes sense why it's significantly higher. Once you remove the first valence electron, aluminum is going to have a net charge of plus one. And so there's going to be less shielding between this electron and the nucleus. So the nucleus is going to have a tighter grip on that electron. Plus there's a net positive charge. And as we know, whenever you have a positively charged ion, the size decreases. So if you decrease the distance between the valence electron and the nucleus, the ionization energy will increase. The nucleus will have a stronger hold on that electron. So the second ionization energy is always higher than the first one. And the third ionization energy will be higher than the second. Once you remove the first two electrons, The aluminum ion now has a charge of plus 2. The energy required to remove the third electron is 2,740, significantly higher than the last one. Now what about the fourth ionization energy? The fourth ionization energy is associated with the removal of an inner core electron, and the energy required to take off one of those electrons is 2,740, is much higher it's 11,600 so as you can see because the inner core electrons are very very close to the nucleus and because they're less shielded from the nucleus the energy that's required to remove such an electron is significantly higher So because aluminum has three valence electrons, the jump in ionization energy occurs or is associated with the fourth electron, because that's a core electron. So make sure you understand the concept of that. So the core electrons are very difficult to remove. They're too close to the nucleus. And so the ionization energy for those electrons are very, very high. An element has the following ionization energies 735, 1445, and 7730. Identify the element and the options to choose from are potassium, magnesium, gallium, silicon, arsenic, and sulfur. Now let's look at the ionization energies. Going from the first to the second, the increase is about 700. But from the second to the third, the increase is over 6000. That means the jump occurs after the second ionization energy. So the third ionization energy is associated with the removal of a core electron, which means the first two are valence electrons. So this element has two valence electrons. So we have to find out which of these elements contain two valence electrons. Sulfur is found in group 6A, so sulfur has six valence electrons. Arsenic is found in group 5, it has five valence electrons. Silicon has four. Gallium has three. Magnesium has two. And potassium has one. So the answer is magnesium, because it contains two valence electrons. So the third electron... That's removed from magnesium is a core electron, which is associated with this number, 7730. Now the next topic of interest is electron affinity. Electron affinity is associated with the energy change that occurs when adding an electron to a gaseous atom. Now when you add electrons to non-metals, that really wants electrons, particularly electronegative nonmetals, they tend to release a lot of energy. And whenever energy is released in a reaction, you have an exothermic reaction. So chlorine has a strong desire for electrons. When it acquires an electron, it turns into chloride, and so it releases a lot of energy. So the energy that's released in a reaction is associated with the element's electron affinity. Ionization energy is the opposite. Ionization energy is the energy required to remove an electron from a gaseous atom, whereas electron affinity is the energy required, or the energy change that occurs when an electron is absorbed by a gaseous atom. It's important to know that the halogens, they are the most exothermic in terms of electron affinity. Because they're so electronegative, they really want electrons. So if you add an electron to, let's say, a gaseous fluorine atom, As it turns into fluoride, it's going to release negative 327.8 kilojoules per mole. So it's highly exothermic. Now the fact that it releases so much energy means that as it acquires that electron, it becomes very stable. So if the addition of an electron to a gaseous atom produces a stable ion, it's going to be highly exothermic. If it produces something unstable, chances are it's not going to be very exothermic, it might be even endothermic. So what are the trends for electron affinity? Generally speaking, keyword generally, electron affinity increases or becomes more exothermic as you go to the right. Now, there's a lot of exceptions. It's helpful to know these numbers. 76, 45, 13, 28. What does that mean? These are group numbers. Group 7 is the most exothermic group. When you add an electron to an element in group 7, the halogens, they will release the most amount of energy. Group 6. next and then it's group four but not five four is more exothermic than five and then it's one in three and two and eight are the least exothermic in fact most elements in two and eight are endothermic you have to put in energy to add an electron Now, it's easy to see why group 7 is the most exothermic, because they're the most electronegative, they only need one electron to complete their octet. But why is group 4 more exothermic than group 5? For example, carbon, which is in group 4, is more exothermic than, let's say, nitrogen, which is found in group 5a. Now, carbon has the configuration that ends in 2p2. For nitrogen, it's 2p3. So, if we add an electron to carbon and one to nitrogen, which one is going to be more stable? Currently, carbon has two electrons in its 2p orbital. Nitrogen has three. So if we wish to create the C-gaseous ion, we just need to add one electron to carbon. And notice that the orbitals, the 2P orbitals, are still unpaired. So because it has this empty space for an extra electron, the C-ion is fairly stable. Now what about adding an electron to an element in group 5a like nitrogen? Once we add it, notice that there's going to be electron repulsion between the two electrons in this orbital. So this is not a good stable arrangement. And it's because of that electron repulsion, that's why elements in group 4A are more exothermic than elements in group 5A. When you add an electron to carbon, you can create, or you will create, a fairly stable ion. But when you add it to nitrogen, the ion won't be as stable due to the electron repulsion. And so that's why group 4A elements are more exothermic than elements in group 5A. more exothermic than group 5 elements in terms of electron affinity. Now what about let's say group 1 versus group 2. Notice the numbers 76, 45, 13, 28. If you add an electron to a group 1 element, it's going to be exothermic. But if you add it to a group 2 element, most of them are endothermic. So let's understand why. Let's use sodium and magnesium as examples. So the configuration for sodium ends in 3s1. magnesium is 3s2. So magnesium has two paired electrons and sodium only has one. Now the next orbital after 3s is 3p. So if we add an electron to sodium, that electron will go in this empty 3s, well, this half-filled 3s orbital. But if we add an electron to magnesium, we have to put that electron out. higher 3p orbital it's more stable to put an electron in the lower energy level than a higher energy level even if there's going to be some sort of electron repulsion So by adding the electron in its half-filled orbital, the electron affinity for sodium is still relatively exothermic. It's much less than 4 and 5, but it's still more exothermic than group 2. Now for group 2 elements, it's endothermic because we're putting an electron in a higher energy level. And so we've got to add energy to put that in. electron there and so that's why it's endothermic for many of the group two elements and the same is true for group eight elements for example let's consider neon neon is in group eight and it ends in 2p6 So the 2p orbital is completely filled. To add an electron, we need to add it to the 3s orbital. And so because we're adding an electron to a higher energy level, it's going to be endothermic. We need to... add energy to the system to add that electron to that atom. So that's why group 2 and group 8 are endothermic with respect to electron affinity. It's because their electrons, the electrons in their... configuration are completely filled. They're completely paired so to add a new electron, you got to put it in a higher energy level. All of the other groups, 7, 6, 4, 5, 1, and 3, you don't have to add an electron to a higher energy level. you simply have to add it to an unfilled or half filled orbital and so adding electrons to half filled or empty orbitals is usually an exothermic process but to put an electron in a higher energy level that's usually an endothermic process so when you're dealing with electron affinity you have to ask yourself if I add this electron will it create a stable ion or an unstable ion is there room for me to put this electron somewhere? Is there a half-filled or an empty-filled orbital to put it somewhere in? If there is, then it's probably going to be exothermic. If you've got to put it in a higher energy level, then chances are it's going to be endothermic. So hopefully, these principles will help you to understand the concept of electron affinity and when it will be endothermic versus when it will be exothermic. Now let's try one more problem. This is going to be the last problem for this video. So let's say if we have the following elements, chlorine, phosphorus, argon, magnesium, sodium, and silicon. Rank the following elements in terms of their electron affinity values. So you want to rank it from endothermic to most exothermic. Now, it helps to know the order. 76, 45, 13, 28. So, as you travel this way towards group 7, it's going to be mostly exothermic. Towards group 2 and 8, it's going to be endothermic. So, chlorine is... in group 7, phosphorus is in group 5, argon is in group 8, magnesium is in group 2, sodium is in group 1, silicon is in group 4. So from endo to exo, let's start with group 8. So argon is going to be endothermic, and magnesium, which is in group 2, that's endothermic. Now we don't have any group 3 elements, so we can get rid of that. The next one is group 1. which is sodium. And then after that, we have group 5, which is phosphorus. And then group 4, which is silicon. We don't have any group 6 elements. The last one is group 7, which is chlorine. So out of the elements listed, if we add an electron, chlorine is going to be the one that's most exothermic. It's going to release a lot of energy. Whereas argon is going to be the least exothermic or the most endothermic. To add an electron to a gaseous atom of argon, we need to add energy to get that going. So that is it for this video. Thanks for watching and have a great day.

That is, as you go to the left, the size of atoms increases, and as you go down... the atoms get bigger. So let's take two examples, hydrogen and helium.

On a periodic table, hydrogen is to the left of helium, so that means that hydrogen is bigger than helium. But now, why is that the case? Because if you look at the atomic mass for helium in the periodic table, it has an atomic mass of 4. And for hydrogen, the atomic mass is 1. So how is it that hydrogen is bigger when helium is more massive? To understand this, we need to jump into physics for a moment.

You know that light charges repel, and opposite charges attract. So if you have a negative charge and a positive charge, these two will feel a force of attraction. Now if you increase the number of charges, let's say if you have two protons and one electron, the force will be two times as large because the charge was increased.

The more protons you have, the greater the force of attraction between the protons and electrons. The second factor to consider is distance. If you have a proton and an electron that are very far apart, the force that they feel will be relatively weak.

But let's say if you cut the distance in half. So let's say currently they're about 4 meters apart. But if you bring the proton and the electron closer, let's say if they're 2 meters apart, If you decrease the distance by a factor of 2, the force of attraction between these two particles will increase by a factor of 4. So the electrostatic traction is much larger if the charges are brought closer to each other.

So if you increase the charge... of any of the protons or electrons, the force of attraction increases. So the more protons that a nucleus has, the greater the force of attraction between the nucleus and the electrons.

And the second factor is distance. If you increase the distance, the force of attraction will decrease. And if you decrease the distance, the electrostatic force of attraction will increase. Now, let's think about what this means. Inner core electrons that are very close to the nucleus, they will feel a very strong force of attraction between itself and the nucleus.

Whereas electrons that are away from the nucleus, like valence electrons, the force of attraction between those electrons and the nucleus will be relatively weak, since they're further away. Now, hydrogen has only one proton in its nucleus, so the charge of the nucleus is plus one, and it has one electron. Now, helium has two protons in its nucleus, and it has two electrons. So because the nuclear charge of helium is greater than that of hydrogen, the force of attraction between those electrons and the nucleus is much larger in helium.

As a result, the size of the... electron cloud shrinks as those electrons are brought closer to the nucleus. And that's one reason helium is smaller than hydrogen even though it has a greater atomic mass. Because the charge of the nucleus increases, the size decreases.

So as you travel from left to right across the periodic table, the atomic size will decrease as the number of protons increases. Now, let's say if we were to compare lithium versus hydrogen. Which atom is bigger?

Now, lithium has three protons and hydrogen has one. yet lithium is bigger than hydrogen. Why is that? I mean, don't lithium have more protons?

Now, the trend as you go from left to right, the size decreases because... because the number of protons increases, but the number of shells is the same. Here, hydrogen and lithium are not in the same row.

They're in different rows. Here's hydrogen, here's helium, and lithium is below hydrogen. Now, for two elements in the same row, the one that's on the right is usually the smaller atom. But now, for two atoms that are in the same group, the one that's below it is usually the bigger one. Remember.

atomic size increases as you go down. So now let's talk about why. Why is it that lithium is bigger than hydrogen? So hydrogen has one proton in its nucleus and it has one electron.

Now lithium has an atomic number three, so it has three protons in its nucleus. Now in the first shell, it can only fit two electrons. These are core electrons. So, it has to put the third electron in the second shell.

And because that electron is further away from the nucleus, the force of attraction between this electron and the nucleus is reduced. Remember, this is the force of attraction. If you increase the distance, the force of attraction decreases. So the reason why lithium is bigger than hydrogen is because it has two electron shells as opposed to one. And that's why as you go down the group in a periodic table, atomic size increases because for every row that you add, you add a new shell of electrons.

And so the atom gets bigger. Now, there's another reason that... you need to consider and that is the concept of shielding.

The inner two core electrons partially shield the outer valence electron from the nucleus. Now this electron is attracted to the three protons in the nucleus. As we said before whenever you have opposite charges they attract each other.

So what if you have like charges let's say two electrons like charges repel each other. So, this electron is repelled by these two electrons. So, even though it feels a force of attraction between... itself in the nucleus, it's also repelled by the other electrons between it and the nucleus.

So you have to consider also the shielding effects of the inner core electrons. Now let's work on some examples. Let's say if we have three elements, chlorine, magnesium, and phosphorus. Rank the following elements in order of increasing atomic size. So that's from small to large.

So which of these atoms is the smallest and which one is the largest? So the first thing you want to do is you want to look at the periodic table. And you want to place them in order. Magnesium is in the third row of the periodic table.

and it's in group 2. It's an alkaline earth metal. Phosphorus is also in the same role and it's in group 5a and chlorine is a halogen in group 7a. Now because these three are in the same row, this is going to be fairly easy. We know that the atomic size increases, or rather it decreases, from left to right.

So that means that magnesium is bigger than phosphorus, and phosphorus is bigger than chlorine. So if we want to rank it in order of increase in atomic size, we need to put the smallest first, and then the largest last. So magnesium is bigger than phosphorus, which is bigger. Let's try another example. Let's say if we have calcium, beryllium, and strontium.

Rank the following elements in order of decreasing atomic radii. So notice that these three elements are all alkaline earth metals. They're found in group 2A. So let's place them in order. So first is BE.

Below BE is magnesium, and then below that is calcium, and then strontium. Now we know that atomic size increases as you travel down a group. Now we want to rank it in order of decrease in atomic size So we know that strontium is the largest Beryllium is the smallest and calcium is in between. So to put it in decreasing order we want to rank it from large to small.

So we're going to start with strontium, then calcium, and then calcium is bigger than beryllium. And so that's how you can rank it in order of decrease in atomic size. Now let's try one more example.

So this time let's say if we have four elements. Fe, Cs, Sulfur, and Helium. Rank the following elements in order of increasing atomic size. So the first thing you want to do is you want to place every element in their respective positions based on the periodic table. Cesium is an alkaline metal in group 1, but it's in the 6th row, so it's like over here.

Fe is a transition metal, and it's in the 4th row, so it's somewhere in this vicinity. Sulfur is a chalcogen found in group 6a, and helium is a noble gas that's in the first row. So remember, atomic size increases as you go to the left and as you go down. So basically... it increases in this general direction, which tells us that cesium is the largest, then it's Fe, then it's S, and then it's helium.

So we wish to rank it in order of increase in atomic size so that's from small to large so helium is smaller than sulfur and sulfur is less than Fe and iron metal is smaller than an atom of cesium so this is the answer now let's move on to ionic radii So consider an atom of lithium and a lithium ion. An atom is a particle that has an equal number of electrons and protons, so atoms are electrically neutral. Ions have unequal number of protons and electrons and so they have a charge. An ion with a positive charge means that there are more protons than electrons. So lithium is an atom with three protons, so the charge of the nucleus is three.

And as we mentioned before, it has two electrons in its first shell, and one valence electron in its second shell. The lithium plus ion, which is also known as a cation, cations are positively charged ions. It has three protons, but it lost an electron. So basically, it lost an entire shell.

So therefore, positively charged cations are significantly smaller than their parent atoms. Because they have less electrons and therefore less electron shells, they will be significantly smaller. Now what about anions, or negatively charged ions?

How do the sizes of these ions compare with their parent atoms? So let's use nitrogen as an example, and the nitride ion. Nitrogen has seven protons in its nucleus. And an atom of nitrogen has seven electrons. So in the first shell, it's going to have two.

And in the second shell, it's going to have 5. So nitrogen has 2 core electrons, or inner electrons, and it has 5 valence electrons. The valence electrons are the electrons in the outermost energy level. Now the nitride ion has 3 more electrons than protons. It has 7 protons and 10 electrons. Protons have a positive charge, electrons have a negative charge.

If you add these two numbers, you're going to get a net charge of minus 3. So let's draw the picture for the nitride ion. Now the first shell contains two electrons, but the second shell has a total of 8 electrons, so that the total number of electrons is 10. So because of the extra three electrons, there's more electron repulsion, which causes the second shell to expand. So negatively charged ions, or anions, are significantly larger than apparent atoms.

So make sure you remember this for ionic radii. Ions with a positive charge will be relatively small, and ions with a negative charge will be relatively large. Now the trend for ionic radii is very similar to that for atomic radii, meaning ionic radii increases as you go down and increases as you go towards the left.

This is especially true for ions with similar charges, if you're comparing two ions with a positive charge or two ions with a negative charge. Let's look at the elements of ions in a row. The sodium ion is bigger than the magnesium ion.

Both of these two ions are isoelectronic. They have the same number of electrons, but because magnesium has more protons than sodium, magnesium is going to be smaller. And aluminum, which has a plus 3 charge, is even smaller than magnesium. So as you can see, as the charges increases, the size of the ion decreases. And as you travel to the right, it decreases to the right, but ionic radii increases towards the left.

Now as you jump from the positive ions towards the negative ions, the size will greatly increase. For example, phosphide, which has a negative 3 charge, it is significantly bigger than almost all of these positively charged ions on the left. And then the size decreases as you go towards the right. So the sulfide ion is smaller than the phosphide ion. And then the chloride ion, which, let me see if I can fit it here, is even smaller than the sulfide ion.

So the size generally decreases as you go from left to right. But if you're comparing a negatively charged ion with a positively charged ion, generally speaking, the negatively charged ion will be bigger than the positively charged ion. Now what about as we go down a group?

Ionic size increases as you go down. For example, the lithium ion is smaller compared to the sodium ion. The sodium ion is larger.

And the potassium ion is even bigger than the sodium ion. So as you go down, the number of electron shells increases, and therefore the ions will increase in size. So if you compare the nitride ion with the phosphide ion, the nitride ion is pretty big, but the phosphide ion is even bigger.

So ionic radii increases as you go down a group. Now let's try a problem. Let's say if we have the following ions. The fluoride ion, the magnesium cation, an atom of neon, the sodium cation, and the oxide ion. Now rank the following atoms and ions in order of increasing size.

Now, to do that, we need to determine if any of these atoms or ions are isoelectronic, meaning that if they have the same number of electrons. If they're isoelectronic, then all you need to do to rank them in order of increase in size is look at the atomic number. For particles that have the same number of electrons, the ones that have more protons will be smaller in size.

As you increase the amount of protons in the nucleus. The increased nuclear charge will cause the electrons to move towards the nucleus, decreasing the size of the atom. So let's count the electrons first. The number of electrons is equal to the atomic number minus the charge.

So if we consider fluorine first, fluorine has an atomic number of 9 and a mass number of 19. So it's 9 minus negative 1. So fluorine has 10 electrons. Now what about magnesium? If you look at the periodic table, magnesium has an atomic number of 12 and a mass number of 24. The small of these two numbers is the atomic number.

So for magnesium it's going to be 12 minus a charge of plus 2 which is 10. So magnesium has 10 electrons. Now for neon it's an atom. which has an atomic number of 10, a mass of 20, and for atoms they're neutral, so they don't have a charge. So it's 10 minus a charge of 0. Neon 2 has 10 electrons. Now what about sodium?

Sodium has an atomic number of 11 and for oxygen has an atomic number of 8. So for sodium it's 11 minus plus 1 which is equal to 10 and for oxygen is 8 minus negative 2. oxygen has 10 electrons. So all of these particles are isoelectronic with each other. They have the same electron configuration. So now all we need to do is look at the atomic number.

atomic number of fluoride is 9, magnesium is 12, neon is 10, sodium is 11, and oxygen is 8. And as you mentioned before, negatively charged ions are generally larger than positively charged ions. And one of the most protons, or the highest atomic number, is going to be the smallest one. So oxide. We want to rank it in order of increase in size, so let's start with the smallest one.

Magnesium has the most or the greatest number of protons. So it's going to be the smallest. It's smaller than sodium and sodium is smaller than neon Which is smaller than fluoride and the oxide ion is going to be the biggest So let's draw the relative sizes.

So this is the smallest. That's a little bit bigger. Then this is going to be bigger. And then the size just continues to increase. So as we can see, oxide has the least amount of oxygen.

number of protons, so that's why it's bigger. Keep in mind, each of these atoms and ions, they have only two shells of electrons. So for particles that have the same number of shells, the one with the most protons will be the smallest. The one with the least number of protons, in this case oxide, is going to be the biggest. So let's compare the magnesium ion and the oxide ion.

Just to get a better understanding of this concept. So we said that magnesium has an atomic number of 12, which means that the charge of the nucleus is 12, since it has 12 protons. And it has a total of 10 electrons. So here's the first shell, and here is the second shell.

Which, I think I could draw a better circle. So the first shell contains two, and the second shell has eight. Now, oxide also contains ten electrons, but the charge is eight, since oxygen only has eight protons in its nucleus.

So the oxide ion is going to be significantly bigger than the magnesium ion. And we can see why it's bigger. Because magnesium has more protons, the electrostatic force between the nucleus and the electrons is stronger because of the increased charge. And so magnesium, it shrinks in size.

The nucleus, it pulls the electrons toward the nucleus, making the ion smaller. Now, in the case of oxygen, the... force of attraction between the outer electrons and the nucleus is relatively weaker because the charge of the nucleus is less. Plus, since the net charge is negative, there's a lot more electron repulsion and so the atom expands when you have extra electron repulsion. So to summarize what we just went over, remember negatively charged ions are usually bigger than positively charged ions.

And for isoelectronic species, that's particles that have the same number of electrons. which is going to have the same number of electron shells, the ones with more protons will be smaller than the ones with less protons. So as you increase the nuclear charge, and if the number of shells remain constant, the particle size will decrease. That extra nuclear charge will cause the electrons to contract towards the nucleus. So now let's move on to electronegativity.

Electronegativity is the ability of an atom to attract an electron to itself. Electronegativity increases towards fluorine. Fluorine is the most electronegative element in the periodic table.

On the upper right corner of the periodic table you have the nonmetals and on the left side you have the metals. Metals tend to be electropositive. They like to give away electrons and form positively charged ions. Nonmetals like to acquire electrons and form negatively charged ions.

So nonmetals tend to be electronegative. Electronegativity increases as you go up and as you travel towards the bottom of the table. towards the right.

So fluorine is more electronegative than oxygen, and oxygen is more electronegative than nitrogen. So let's say if you have a fluorine atom, and if you add an electron to it, fluorine will turn into the fluoride ion. The nonmetal will become a negatively charged anion.

The nonmetals tend to be electronegative. they like to acquire electrons. Now if you have an atom of sodium, sodium wants to give away an electron to form a positively charged cation. So metals tend to be electropositive. They like to give away electrons.

So those are some things to know. Now let's say if you have four elements, silicon, magnesium, chlorine, and aluminum. Rank the following elements in order of increasing electronegativity.

So, look at the periodic table and place them in order. So, each of these elements are in the same row. Magnesium comes first, it's in group 2. Then aluminum, which is in group 3A.

And then silicon, that's in 4A. And chlorine is in group 7A. So electronegativity increases towards the right.

So if we want to rank it in order of increasing electronegativity, that's from low to high. Magnesium is going to be the least electronegative, it's a metal. And then it's aluminum, which is also a metal.

And then silicon, which is a metalloid. And then chlorine is the most electronegative, that's a nonmetal. So nonmetals are usually...

So this is the answer. Now what about these? Let's say if we have tin, germanium, lead, and carbon. Rank the following elements in order of decreasing electronegativity.

Now each of these elements are found in group 4A of the periodic table. So let's put them in order. First we have carbon, and then below that, GE. germanium and then tin metal and then lead.

Electronegativity increases as you travel up a group so we want to rank it in decreasing order so we need to put the most electronegative element. We need to rank it from high to low. So the highest is carbon which is a nonmetal, then is germanium which is a metalloid, and then tin metal, and then lead metal. So the metals, they have the lowest electronegativity.

So this is the answer. That's how you can rank it in order of decreasing electronegativity. Now, which element is more electronegative, silicon or nitrogen?

If you place these elements in their respective position, silicon is in group 4A in the third row, nitrogen is in group 5A. in the second row. Electronegativity increases towards fluorine. So it increases as you go up and to the right.

So therefore you can clearly see that nitrogen is more electronegative than silicon. Now what about nitrogen versus sulfur. If you place them in their respective positions, nitrogen is in group 5A in the second row, sulfur is in group 6A but in the third row.

But electronegativity increases this way. So which one is more electronegative? So as you travel up, electronegativity increases, but as you travel to the right, it decreases.

So sulfur is to to the right of nitrogen, but it's also below nitrogen. So how can you tell which one is more electronegative in this case? It turns out that the electronegativity increases more when you go up than when you go to the left.

So the increase for going up one row is greater than a decrease from traveling one unit to the right. Now if you're ever unsure, you can always use the equation that I just showed you. So the increase for going up one row At this point, you can look at the election negativity table and see which one is higher. The election negativity for nitrogen, according to most textbooks, is roughly about 3.0.

For sulfur, it's 2.5. So as you can see, going up has more priority than going to the right. So let's go over some common electronegativity values for elements like boron, carbon, nitrogen, oxygen, fluorine, phosphorus, sulfur, chlorine, bromine, and iodine.

The election negativity value for boron is about 2.0. For carbon, it's 2.5. For anode, it's 3.0. For O, it's 3.5. And for fluorine, it's about 4.0.

Now for phosphorus, it's roughly around 2.0. one. For sulfur 2.5 and for chlorine it's about 3.0. For bromine 2.8, iodine 2.5. So as you can see there's a large increase as you go from to rho 3. Here the increase is 0.9 and here it's about 1. So that's why nitrogen was significantly higher than that for sulfur.

So if you need to go this way, generally speaking, this is the one that's going to win. But not always though. For example, bromine is higher than sulfur.

So there are some exceptions. So you may just need to know some of these values if you ever get an unusual question like that. By the way, the electronegativity for hydrogen is about 2.1. So even though hydrogen is to the left of boron, boron is in row 2, hydrogen is in row 1. So hydrogen is a little bit more electronegative than boron.

Now the next train that we need to talk about is metallic character metallic character increases in this general direction Towards the metals the nonmetals are located in the upper right corner of the periodic table and the metals is just to the left But metallic character increases as you travels to the left and down across the periodic table So let's go over some examples. So let's say if we have elements such as silicon sodium sulfur, aluminum, and chlorine. Rank the following elements in order of increase in metallic character.

So as we've been doing before let's place it based on their respective positions in a periodic table. Now each of these elements are located in the same row in the periodic table. Sodium is found in group 1 and then aluminum in group 3a.

silicon in group 4a, sulfur in group 6a, and then Cl in group 7a. So metallic character increases as you travel towards the left on the periodic table. So if we wish to rank it in order of, let's say, increasing metallic character, we need to rank it from low to high. So chlorine, which is a nonmetal, is going to have the lowest metallic character. Sulfur is also a nonmetal.

Silicon is a metalloid, aluminum is a metal, and sodium is a metal. Now let's think about what this means. Sodium really wants to get rid of its electrons more than aluminum. Both of these elements are metals, but sodium has a greater metallic character than aluminum, so sodium is more electropositive.

It wants to get rid of its electrons with a stronger force than aluminum. Now between sulfur and chlorine, chlorine is more electronegative, so it's less metallic than sulfur. Silicon is a metalloid.

It's in between a metal and a nonmetal. So this is the answer. That's how you can rank it in order of increasing metallic character.

Now metallic character increases as you go down a group. Let's consider the group 4A elements, like carbon. silicon, germanium, tin, and lead. As you go down the periodic table, the elements change from nonmetals to metals.

Carbon is a nonmetal, silicon is a metalloid, and germanium is a metalloid. But tin is a metal and lead is a metal. So as you can see, metallic character increases as you travel down a group. Let's try one more example.

Let's say if we have Gallium, Manganese, Nitrogen, Helium, and Francium. Rank the following elements in order of decreasing metallic character. Francium is located in group 1 in the 7th row.

And then you have manganese, which is a transition element. And then gallium is to the right of that. And then there's nitrogen and helium. So metallic character increases in this general direction. So if we want to rank it and decrease in order, we need to rank it from high to low.

So francium has the greatest metallic character. Then it's manganese, and then Ga, and then N. and then He.

The next topic that we need to talk about is ionization energy. Ionization energy is the energy required to remove an electron from a gaseous atom. Now generally speaking, it's easier to remove an electron from a metal than it is to remove from a nonmetal.

Metals like to give away electrons. So it doesn't require that much energy to remove an electron from it. Nonmetals like to acquire electrons, so it's harder to remove an electron from a nonmetal. The trend for ionization energy is in this direction.

Ionization energy increases towards helium. It increases as you go up across the periodic table and towards the right. So the main reason why it increases as you travel from left to right... is due to the increase in nuclear charge. From left to right, the number of electron shells, or the principal quantum number, remains the same.

However, the number of protons in the nucleus increases. And as you add more protons to the nucleus, those protons will have a tighter grip on the electrons. So it requires more electrons, I mean, it requires more energy to remove those electrons.

So let's compare lithium and beryllium. Lithium has three protons in its nucleus, and beryllium has a nuclear charge of four. Now keep in mind, beryllium is smaller than lithium because beryllium has a higher nuclear charge. So here's a question for you.

Is it easier to remove the electron from lithium or a valence electron from beryllium? It's going to be easier to remove the electron from lithium. The ionization energy for lithium is about 520, but for beryllium it's about 899. So it requires much more energy to remove an electron from beryllium than to remove it from lithium.

For one thing... The force of attraction between the electron and the nucleus is weaker compared to the force of attraction between an electron and beryllium than the nucleus. The nuclear charge of beryllium is much greater than that of lithium, so the nucleus has a tighter grip on this electron.

So that's one reason why it's harder to remove it. But also, beryllium is smaller than lithium, and so the distance between the electron and the nucleus is much less than the valence electron and the nucleus within lithium. So remember, if you decrease the distance, the force of attraction increases.

So because beryllium is smaller, and plus the fact that it has a higher nuclear charge, means that beryllium... holds on to that valence electron with a tighter grip than lithium. And so that's why it requires so much more energy to remove that valence electron in beryllium. It has to do with the increased nuclear charge and the smaller size.

Both of those factors increases the ionization energy of BE. Now what about between lithium and sodium? Which element has a higher ionization energy?

Sodium is below lithium and ionization energy increases as you go up. The ionization energy for sodium is about 495, but for lithium it's 520, so it doesn't vary much. Now lithium, as we mentioned before, has a nuclear charge of 3, and sodium has 11 protons. So sodium is bigger than lithium. It has three electron shells.

So why is it easier to remove a valence electron from sodium than it is to remove it from lithium, considering that sodium has a higher nuclear charge? Now granted, as the charge increases, we know that the electrostatic force of attraction increases. So an atom with a higher nuclear charge means that it has electrons that are harder to remove. However, you have to take into account distance.

The distance between this electron and the nucleus is significantly larger than the distance between this electron and the nucleus. So if you increase the distance, the force greatly decreases. And so because this electron is very far away from the nucleus, it's relatively easy to remove that electron. And that's why sodium has a slightly lower ionization energy than lithium.

It's because of the increased distance between that electron and the nucleus, even though it has a higher nuclear charge. Now we can see that distance plays a greater role in affecting the force of attraction than the charge. Because the electron is further away, the ionization energy is less, even though the nuclear charge is greater. And this is consistent with Coulomb's Law. His law describes the relationship between the electrostatic force between two charged particles.

So let's say if you have a proton and if you have an electron. The force of attraction between these two particles is called the force of attraction. particles is proportional to the charge and inversely related to the distance between them which is R so if you double the value of Q the force will double in value however If you double the distance, because it's squared, 2 squared is 4, the force will reduce by a factor of 4. So the distance between the charges has a greater impact than...

the magnitude of the charge based on the equation. So to summarize what we've learned, as we travel down a group, the ionization energy decreases because the distance between the protons and the valence electrons increases. And as we travel to the right, ionization energy increases due to the increase in the effective nuclear charge.

As you go to the right, the principal quantum number remains the same. So the number of electron shells is roughly about the same. Now there are some exceptions that you need to be aware of.

For example, between beryllium and boron. The ionization energy for boron is 800 and for beryllium the first ionization energy is 899. Now boron is to the right of beryllium in the periodic table so typically we should expect that boron should have a higher ionization energy but it doesn't. The question is why?

Now the last electron in beryllium is the 2s2 electron. However, the last electron in boron is the 2p1 electron. 2p, on average, is farther away from the nucleus than 2s. If you write the electron configuration, it's 1s2, 2s2, and then 2p6.

But if you draw the orbital diagram energy levels, it looks like this. So 2p is higher in energy, which means that it's further away from the nucleus. And as we mentioned before, if you increase the distance between the electrons and the nucleus, on average, the ionization energy tends to decrease. And so... that's why we see a slight drop as you go from be to boron the same is true from magnesium to aluminum as you go from the s2 level to the the first P level there's a temporary decrease even for calcium to gallium so from magnesium to aluminum it changes from 735 to 580 which is pretty significant From calcium to gallium from group 2a to group 3a It drops from 590 to 570 so I mean 579 so the change between these two is not that great, but These are exceptions that you want to be aware of if you're tested on it Now, there are some other exceptions that you need to be aware of as well.

It's not over yet. And that's going from nitrogen to oxygen, or phosphorus to sulfur, or arsenic to selenium. The first ionization energy for nitrogen is oxygen.

is 1402 and for oxygen is 1314 at least according to the textbook I'm using now why does it temporarily decrease from n to O the electron configuration for nitrogen ends in 2p3 and for oxygen it ends in 2p4 Now for nitrogen, it has three unpaired electrons. But for oxygen, it has a paired electron. This paired electron is the one that's being removed.

And due to the electron repulsion between the two electrons in this orbital, that allows this electron, the one being removed, It allows it to be removed with ease. Because that electron is being repelled by the other electron in that orbital, it doesn't take much energy to remove that valence electron. And that's why we see this drop in ionization energy.

It's due to the electron repulsion between these two paired electrons. Now let's try some examples. Rank the following elements in order of increase in ionization energy.

That is first ionization energy. So let's say if we have the elements gallium, bromine, potassium, chromium, and arsenic. So once again, you want to look at the periodic table, and you want to place them in their respective positions.

So each of these elements are in the first, not the first row, but the fourth row. The first one is potassium, and then after potassium it's chromium, which is a transition metal. And then you have gallium, and then arsenic, and then bromine, which is a nonmetal.

The first ionization energy increases as you travel left to right on a periodic table. So let's say if we want to rank it in order of increasing ionization energy, we could simply write it like this. So this is the answer. Now what about these three? Phosphorus.

arsenic, nitrogen, and antimony. Go ahead and rank these elements in order of increasing or rather decrease in first ionization energy. Now these elements are all found in group 5a of the periodic table.

So if we place them in order, first we have nitrogen, below that is phosphorus, then arsenic, and then after that, antimony. Now, the first ionization energy increases as you go up a group within the periodic table. So, if we want to rank it and decrease in order from high to low, we need to start with the highest, which is nitrogen.

And that's greater than phosphorus, which is larger than arsenic, and which has a higher first ionization energy than antimony. So, that's how you can rank it in order of decrease in first ionization energy. Now let's try another example.

Let's say if we have fluorine, phosphorus, helium, francium, and vanadium. So let's put this in order. So francium is all the way at the bottom, towards the left. It's in the 7th row, 1st column.

Vanadium is a transition metal which is in this region in the periodic table. And then we have phosphorus, fluorine, and helium is to the upper right of fluorine. So first ionization energy increases towards helium. So therefore, let's say if we want to rank it in order of increase in ionization energy from low to high, let's start with francium, which is less than vanadium, that's less than P, which is less than F, and that's less than He.

So it helps if you can place the elements based on where they're located in a periodic table, and then use the trend to rank them in the appropriate order. Now sometimes you may have questions about the second ionization energy and the third and so forth. The second ionization energy is the energy required to remove the second electron. That second electron may be a valence electron or maybe a core electron. So let's consider aluminum.

Aluminum is in group 3a and it has a total of 13 electrons. Its atomic number is 13 so as an atom A neutral atom has 13 protons and 13 electrons. Now let's draw the aluminum atom. So the charge of the nucleus is 13. In the first shell, it's going to have 2 electrons. In the second energy level, it's going to have 8 electrons.

So right now we have a total of 10. And in the third energy level, it has three valence electrons. Now the first ionization energy of aluminum is the energy required to remove the first valence electron. Let's say this one. So it's going to take about 580 kilojoules of energy to remove one mole of electrons from aluminum. And that's really one electron per one aluminum atom, but it's 580 kilojoules per mole.

Now, for the second electron, the first ionization energy is higher. It's 1815. And it makes sense why it's significantly higher. Once you remove the first valence electron, aluminum is going to have a net charge of plus one. And so there's going to be less shielding between this electron and the nucleus.

So the nucleus is going to have a tighter grip on that electron. Plus there's a net positive charge. And as we know, whenever you have a positively charged ion, the size decreases.

So if you decrease the distance between the valence electron and the nucleus, the ionization energy will increase. The nucleus will have a stronger hold on that electron. So the second ionization energy is always higher than the first one.

And the third ionization energy will be higher than the second. Once you remove the first two electrons, The aluminum ion now has a charge of plus 2. The energy required to remove the third electron is 2,740, significantly higher than the last one. Now what about the fourth ionization energy?

The fourth ionization energy is associated with the removal of an inner core electron, and the energy required to take off one of those electrons is 2,740, is much higher it's 11,600 so as you can see because the inner core electrons are very very close to the nucleus and because they're less shielded from the nucleus the energy that's required to remove such an electron is significantly higher So because aluminum has three valence electrons, the jump in ionization energy occurs or is associated with the fourth electron, because that's a core electron. So make sure you understand the concept of that. So the core electrons are very difficult to remove. They're too close to the nucleus.

And so the ionization energy for those electrons are very, very high. An element has the following ionization energies 735, 1445, and 7730. Identify the element and the options to choose from are potassium, magnesium, gallium, silicon, arsenic, and sulfur. Now let's look at the ionization energies.

Going from the first to the second, the increase is about 700. But from the second to the third, the increase is over 6000. That means the jump occurs after the second ionization energy. So the third ionization energy is associated with the removal of a core electron, which means the first two are valence electrons. So this element has two valence electrons.

So we have to find out which of these elements contain two valence electrons. Sulfur is found in group 6A, so sulfur has six valence electrons. Arsenic is found in group 5, it has five valence electrons. Silicon has four.

Gallium has three. Magnesium has two. And potassium has one. So the answer is magnesium, because it contains two valence electrons. So the third electron...

That's removed from magnesium is a core electron, which is associated with this number, 7730. Now the next topic of interest is electron affinity. Electron affinity is associated with the energy change that occurs when adding an electron to a gaseous atom. Now when you add electrons to non-metals, that really wants electrons, particularly electronegative nonmetals, they tend to release a lot of energy.

And whenever energy is released in a reaction, you have an exothermic reaction. So chlorine has a strong desire for electrons. When it acquires an electron, it turns into chloride, and so it releases a lot of energy. So the energy that's released in a reaction is associated with the element's electron affinity. Ionization energy is the opposite.

Ionization energy is the energy required to remove an electron from a gaseous atom, whereas electron affinity is the energy required, or the energy change that occurs when an electron is absorbed by a gaseous atom. It's important to know that the halogens, they are the most exothermic in terms of electron affinity. Because they're so electronegative, they really want electrons.

So if you add an electron to, let's say, a gaseous fluorine atom, As it turns into fluoride, it's going to release negative 327.8 kilojoules per mole. So it's highly exothermic. Now the fact that it releases so much energy means that as it acquires that electron, it becomes very stable. So if the addition of an electron to a gaseous atom produces a stable ion, it's going to be highly exothermic. If it produces something unstable, chances are it's not going to be very exothermic, it might be even endothermic.

So what are the trends for electron affinity? Generally speaking, keyword generally, electron affinity increases or becomes more exothermic as you go to the right. Now, there's a lot of exceptions.

It's helpful to know these numbers. 76, 45, 13, 28. What does that mean? These are group numbers. Group 7 is the most exothermic group. When you add an electron to an element in group 7, the halogens, they will release the most amount of energy.

Group 6. next and then it's group four but not five four is more exothermic than five and then it's one in three and two and eight are the least exothermic in fact most elements in two and eight are endothermic you have to put in energy to add an electron Now, it's easy to see why group 7 is the most exothermic, because they're the most electronegative, they only need one electron to complete their octet. But why is group 4 more exothermic than group 5? For example, carbon, which is in group 4, is more exothermic than, let's say, nitrogen, which is found in group 5a. Now, carbon has the configuration that ends in 2p2. For nitrogen, it's 2p3.

So, if we add an electron to carbon and one to nitrogen, which one is going to be more stable? Currently, carbon has two electrons in its 2p orbital. Nitrogen has three.

So if we wish to create the C-gaseous ion, we just need to add one electron to carbon. And notice that the orbitals, the 2P orbitals, are still unpaired. So because it has this empty space for an extra electron, the C-ion is fairly stable.

Now what about adding an electron to an element in group 5a like nitrogen? Once we add it, notice that there's going to be electron repulsion between the two electrons in this orbital. So this is not a good stable arrangement.

And it's because of that electron repulsion, that's why elements in group 4A are more exothermic than elements in group 5A. When you add an electron to carbon, you can create, or you will create, a fairly stable ion. But when you add it to nitrogen, the ion won't be as stable due to the electron repulsion.

And so that's why group 4A elements are more exothermic than elements in group 5A. more exothermic than group 5 elements in terms of electron affinity. Now what about let's say group 1 versus group 2. Notice the numbers 76, 45, 13, 28. If you add an electron to a group 1 element, it's going to be exothermic.

But if you add it to a group 2 element, most of them are endothermic. So let's understand why. Let's use sodium and magnesium as examples. So the configuration for sodium ends in 3s1. magnesium is 3s2.

So magnesium has two paired electrons and sodium only has one. Now the next orbital after 3s is 3p. So if we add an electron to sodium, that electron will go in this empty 3s, well, this half-filled 3s orbital. But if we add an electron to magnesium, we have to put that electron out. higher 3p orbital it's more stable to put an electron in the lower energy level than a higher energy level even if there's going to be some sort of electron repulsion So by adding the electron in its half-filled orbital, the electron affinity for sodium is still relatively exothermic.

It's much less than 4 and 5, but it's still more exothermic than group 2. Now for group 2 elements, it's endothermic because we're putting an electron in a higher energy level. And so we've got to add energy to put that in. electron there and so that's why it's endothermic for many of the group two elements and the same is true for group eight elements for example let's consider neon neon is in group eight and it ends in 2p6 So the 2p orbital is completely filled.

To add an electron, we need to add it to the 3s orbital. And so because we're adding an electron to a higher energy level, it's going to be endothermic. We need to...

add energy to the system to add that electron to that atom. So that's why group 2 and group 8 are endothermic with respect to electron affinity. It's because their electrons, the electrons in their... configuration are completely filled. They're completely paired so to add a new electron, you got to put it in a higher energy level.

All of the other groups, 7, 6, 4, 5, 1, and 3, you don't have to add an electron to a higher energy level. you simply have to add it to an unfilled or half filled orbital and so adding electrons to half filled or empty orbitals is usually an exothermic process but to put an electron in a higher energy level that's usually an endothermic process so when you're dealing with electron affinity you have to ask yourself if I add this electron will it create a stable ion or an unstable ion is there room for me to put this electron somewhere? Is there a half-filled or an empty-filled orbital to put it somewhere in? If there is, then it's probably going to be exothermic. If you've got to put it in a higher energy level, then chances are it's going to be endothermic.

So hopefully, these principles will help you to understand the concept of electron affinity and when it will be endothermic versus when it will be exothermic. Now let's try one more problem. This is going to be the last problem for this video. So let's say if we have the following elements, chlorine, phosphorus, argon, magnesium, sodium, and silicon.

Rank the following elements in terms of their electron affinity values. So you want to rank it from endothermic to most exothermic. Now, it helps to know the order.

76, 45, 13, 28. So, as you travel this way towards group 7, it's going to be mostly exothermic. Towards group 2 and 8, it's going to be endothermic. So, chlorine is... in group 7, phosphorus is in group 5, argon is in group 8, magnesium is in group 2, sodium is in group 1, silicon is in group 4. So from endo to exo, let's start with group 8. So argon is going to be endothermic, and magnesium, which is in group 2, that's endothermic.

Now we don't have any group 3 elements, so we can get rid of that. The next one is group 1. which is sodium. And then after that, we have group 5, which is phosphorus.

And then group 4, which is silicon. We don't have any group 6 elements. The last one is group 7, which is chlorine.

So out of the elements listed, if we add an electron, chlorine is going to be the one that's most exothermic. It's going to release a lot of energy. Whereas argon is going to be the least exothermic or the most endothermic. To add an electron to a gaseous atom of argon, we need to add energy to get that going.

So that is it for this video. Thanks for watching and have a great day.

Transcript for:Understanding Periodic Trends in Chemistry

Transcript for:
Understanding Periodic Trends in Chemistry