Chapter 7 - Periodic Properties of the Elements

Why do elements react more dramatically than others? If we grew up a piece of gold metal into water, nothing happens. However, dropping lithium metal into water initiates a slow reaction in which bubbles readily form on the surface of the metal. By contrast, potassium metal reacts suddenly and violently with water. Why do lithium and potassium react so differently with water, even though they are from the same family of the periodic table? To understand such differences, we will examine how some atomic properties change systematically as we traverse the periodic table. Here in Chapter seven, by the end of this chapter, you will be able to describe the organization of the periodic table and explain its significance. You will also use the periodic table to identify the common ions formed by various elements, use the concept of effective nuclear charge to explain the basis for observed periodic trends, identify periodic trends in ionization energies, electron affinities, and atomic radii, and explain how the radius of an atom changes upon losing electrons to form a cation or gaining electrons to form an ion. Furthermore, you will be able to write the electron configuration of ions. The discovery of chemical elements has been ongoing since ancient times. Certain elements, such as gold, appear in nature in elemental form and were discovered thousands of years ago. In contrast, some elements, like technetium, are artificially created and intrinsically unstable. We know about them only because of technology developed during the 20th century. The majority of elements primarily form compounds and consequently are not found in nature in their elemental form. For centuries, scientists were unaware of their existence. During the early to mid-19th century, advancements in chemistry made it easier to isolate elements from their compounds. As a result, the number of known elements more than doubled, from 31 in the 1800s to 63 by 1865. As the number of known elements increased, scientists began classifying them. In 1869, Dmitri Mendeleev in Russia and Lothar Meyer in Germany published nearly identical classification schemes. Both noted that similar chemical and physical properties recur periodically when the elements are arranged in order of increasing atomic weight. At this time, scientists had no knowledge of atomic number, but they knew that atomic weight generally increased with increasing atomic number. So, Mendeleev and Meyer fortuitously arranged the elements in nearly the proper sequences. Mendeleev is given more credit for advancing this idea more vigorously, as his simultaneous work was more well-known. For example, here was one of Mendeleev's predictions for an element that, at his time, didn't exist, but he expected it to be somewhere. He predicted that the atomic weight of this element would be 72, the density 5.5, the specific heat 0.305, and most importantly, its melting point would be high. The color of this element would be dark gray, and the formula for its oxide would be XO2, with a density of 4.7. The formula for the chloride of this element would be XCl4, and its boiling point would be under 100 degrees. In 1871, when they discovered the element germanium, they found that its atomic weight was 72.59, which was close to Mendeleev's prediction. The density was 5.35, the specific heat 0.309, and the melting point 947 degrees, which is relatively high. The color was grayish white, and the formula for the oxide was GeO2, similar to what Mendeleev predicted to be XL2. The density of the oxide was close to what he expected as well. The formula for eyford mining is GeCl4, and he observed that it was basically predicted as XeCl4. The boiling point for that chloride was 84 degrees, expected to be under 100. So, most of the credit was given to Mandala because he was more aggressive in his prediction. To share all this knowledge, Mandala's table was based on atomic masses, the most fundamental property of elements known at that time. About 35 years later, the nuclear atom was discovered by Ernest Rutherford, and two years later, scientists had mostly developed the concept of atomic number experimentally. The number of protons was considered the basis for the periodic property of the elements, and that's how the atomic number was established. Basically, all this combination of knowledge leads to establishing the atomic number, which is used today to organize the periodic table. Periodicity is the repetitive pattern of a property for elements based on atomic number, and that's how the periodic table was organized. The following properties are discussed in this chapter: the size of atoms and ions, ionization energy, electron affinity, and some group of chemical properties trends. First, we will discuss a fundamental property that helps make these trends effective: nuclear charge. The effective nuclear charge is the property that helps with all these trends. Many properties depend on the attraction between the valence electrons and the nucleus. Remember that valence electrons are those electrons in the outermost shell of an atom. So, the valence electrons are both attracted to the nucleus and repelled by the electrons. The forces an electron experiences depend on both factors. The effective nuclear charge (Zeff) is basically the difference between the attraction it feels from the nucleus and the repulsion it feels with other electrons. Zeff is found this way: Zeff = Z - s, where Z is the atomic number and s is a screening constant, usually close to the number of inner electrons. The effective nuclear charge is a periodic property. It increases across a period and decreases down a group. From left to right in the periodic table, the effective nuclear charge increases, and from top to bottom in a group, it decreases. Here we have a figure to explain the effective nuclear charge better. The figure represents a nucleus and an electron. The blue line represents the force of attraction from the nucleus, and the red line represents the repulsion the electron feels from other electrons in the core. The combination of these forces is known as the effective nuclear charge. Now, let's discuss the size of an atom. In a collision between two atoms, the shortest distance separating the two nuclei is twice the radius of the atom. We call this the nonbonding atomic radius or the van der Waals radius. The nonbonding atomic radius is also known as the van der Waals radius. For any atom in a molecule, the atomic radius is equal to half of the bond distance (D). This collision Basically, we create a bond between two atoms, and the non-bonding atomic radius of an atom is half of the bonding atomic radius. This allows us to define the size of an atom. The trend of the size of atoms in the periodic table is that the bonding atomic radius tends to decrease from left to right across a period. This is because the effective nuclear charge, or the positive charge in the nucleus, increases as the atomic number increases from left to right. However, when going from top to bottom in the periodic table, the size of the atom increases. This is because each level, or energy shell, from 1 to 2, for example, is larger, leading to an increase in the size of the atom from top to bottom. The decrease in size from left to right is due to the effect of the increasing effective nuclear charge on the valence electrons. Now, let's consider what happens when ions are formed. The size of ions is determined by the inter-atomic distance in an ionic compound, which is similar to how the size of an atom is determined. In the periodic table, metals produce cations by losing electrons, while nonmetals on the right side of the table produce anions by gaining electrons. The size of ions depends on the nuclear charge, number of electrons, and orbitals in which the electron resides. When ions are formed, cations are smaller than their parent atoms because the outermost electron is removed, reducing electron repulsion. The cations have a larger effective nuclear charge and a reduced radius. On the other hand, anions are larger than their parent atoms because an electron is added, increasing electron repulsion and expanding the electron cloud. The anions have a larger radius. To make it clearer, let's talk about the size of ions in an isoelectronic series. An isoelectronic series consists of ions that have the same number of electrons. For example, in the series oxygen (O^2-), fluorine (F^-), sodium (Na^+), magnesium (Mg^2+), and aluminum (Al^3+), they all have 10 electrons. So that's why oxygen is -2, fluorine is 9, so we add one more and it's at 10. I'm sodium has 11, so we need to remove one in a way that we can have ten electrons in sodium. So that's why the ion is +1. Magnesium is 12, so we need to remove two electrons. So in that way, it is magnesium +2 and we have 10 electrons here. Aluminum is 13, so in a way that we can have 10 electrons, we need to remove three electrons from aluminum. So in this way, all of these ions have just 10 electrons, so they are basically in a nice electronic series and the ionic size decreases with an increasing nuclear charge. Okay, we can see here that as the nuclear effective charge is decreasing, basically it's increasing here because we're removing electrons in a way that the nuclear charge is more positive and that we pull more the shells to the nucleus. We can see as the nuclear charge is basically increasing to the right, the ionic radius is decreasing. Okay, so in this way, we are also showing that in a nice electronic series, atoms that have the same number of electrons, the one that has the higher nuclear charge or effective nuclear charge basically is the one that has the smallest ionic radius. We have another concept that is known as ionization energy. The ionization energy is the minimum energy required to remove an electron from the ground state of a gaseous atom or ion. The first ionization energy is the energy required to remove the first electron. So we have for example one element, one atom is neutral, it has the same number of electrons and the same number of protons. So the energy required to remove that first electron is known as the first ionization energy. The second ionization energy is the energy required to remove the second electron. So now you have basically an ion and you're going to remove a second electron. That's the second ionization energy. These energies are different. The higher the ionization energy, the more difficult it is to remove an electron. That means that not all the elements are going to be easy to remove an electron. Some of them are going to be harder than others. The higher the ionization energy, the more difficult it is to remove an electron. It requires more energy to remove each successive electron. For the third or fourth ionization energy, to remove other ions. You can see here in this table in 7.2, we have here the first ionization energy for the first electron is 496 kilojoules per mole, but to remove the second is 4562 kilojoules per mole. A lot more to remove this electron, the second one. Here in magnesium, it is 700 kilojoules, and for the second one it is 1451. But for the third, to try to remove the first electron, it is 7733 kilojoules per mole. So basically, these are a lot harder to remove. In this case, that's why most of the time studying is just positive 1 ions, positive 2 is really hard to remove this third electron from the magnesium. The same thing with aluminum, to remove this fourth electron is going to be pretty hard. We need to add a lot of energy to remove that electron. But the important thing here is that you can see that for the first one, because compared to the second one, it is easier. It has a lot less energy required to remove that first electron, as for the second, as for the third. For the second, you need a lot more energy to remove the third one because you have a more positive charge in that ion, and to remove an electron is harder. And in that way, as you remove an electron, you are increasing that effective nuclear charge and you are making those electrons closer to the nucleus. So that's why it's really hard to remove the second or even the third or fourth one in some of those ions. Let's see the periodic trend of the first ionization energy. The first ionization energy generally increases across the period and generally decreases down a group, the s and the P. Block elements show a larger range of values for the first ionization energy, while the d-block elements generally increase slowly across the period. The f-block elements show only small variation. So, as you go from left to right in a period, you are increasing the first ionization energy. This is because of the higher effective nuclear charge, which makes it more difficult to remove an electron. On the other hand, when you go from top to bottom, the ionization energy decreases. This is because the electrons are farther from the nucleus and easier to remove in the elements at the bottom of the periodic table. Smaller atoms have higher ionization energy values. As you can see, as you go from left to right, the ionization energy increases due to the increasing effective nuclear charge. And as you go from top to bottom, the ionization energy decreases due to the larger radius and the electrons being farther from the nucleus. Now, let's talk about the electron configuration of ions and cations. When producing a cation, electrons are lost from the highest energy level. For example, the electronic configuration of lithium as an atom is 1s2 2s1. When we produce the cation, we need to remove one electron from the highest energy level, which is 2s1. So the electron configuration for the cation is 1s2. Now, let's consider iron. The electronic configuration for iron as an atom is [Ar] 3d6 4s2. One of the ions that can be produced by iron is Fe+2. In order to create this cation, we need to remove two electrons from the highest energy level, which is 4s2. So the electron configuration for Fe+2 will be [Ar] 3d6. For anions, we need to create the electron configuration of ns2 np6. Let's take the example of fluorine, which as an atom has the electronic configuration 1s2 2s2 2p5. When fluorine gains an electron, its electron configuration becomes 1s2 2s2 2p6, resulting in a negative charge. 2P^6 is the electronic configuration for nitrogen. Another characteristic found in the periodic table is electronic affinity. Electronic affinity is the energy change associated with the addition of an electron to a gaseous atom. It represents the energy associated when an atom gains an electron. For example, when a chlorine atom gains an electron, it forms a chloride ion and releases energy. This reaction is typically exothermic, resulting in a negative electronic affinity. In general, there is not much change in electronic affinity within a group, but it increases across a period. However, there are three notable exceptions: group 2A, group 5A, and group 8. Group 2A and group 8 have full sublevels, while group 5A has a half-filled P sublevel. Now, let's discuss the characteristics of metals, nonmetals, and metalloids. Metals are found on the left side of the periodic table, nonmetals on the right side, and metalloids in between. One difference between metals and nonmetals is their ability to form cations and anions. Metals tend to lose electrons and form cations, while nonmetals tend to gain electrons and form anions. This is because metals aim to achieve the electron configuration of a noble gas by losing electrons, while nonmetals aim to achieve the noble gas configuration by gaining electrons. To elaborate, for example, sodium needs to release one electron to attain the electron configuration of a noble gas, so it easily forms a cation. Similarly, elements in group 2A such as magnesium and calcium can lose two electrons to achieve the noble gas configuration. On the other hand, nonmetals such as those in family 7, group 6, and group 5 can gain electrons to reach the noble gas configuration. This is why nonmetals tend to form anions more easily. Now let's discuss some characteristics of metals. Most elements in nature are metals. They have properties such as shine (luster), good conductivity of heat and electricity, malleability, and ductility. They are solid at room temperature, except for mercury. Metals also have low ionization energy, making it easy for them to form cations by losing electrons. The chemistry of compounds formed between metals and nonmetals tends to be ionic in nature. Note: The paragraphing has been adjusted for better readability. You see a metal in a compound, you can think that the compound is an ionic compound. The metal oxides tend to be basic. Here, we have nitric acid and water, and we also have nickel oxide and nitric acid. As we can see here, when it reacts with water, nickel oxide is insoluble, but it reacts with nitric acid. This means that because it reacts with an acid, the metal is considered a base. Most metal oxides tend to be basic. The nonmetals are found on the right-hand side of the periodic table. The properties of nonmetals include the following: they can be solid, liquid, or gas, depending on the element. The solids are dull, brittle, and poor conductors. They have largely negative electronegativity and can form anions easily due to gaining electrons. The chemistry of nonmetals is largely molecular compounds when they are found in substances containing only nonmetals. Most nonmetal oxides are acidic, unlike metal oxides. The metalloids, on the other hand, have some characteristics of metals and some of nonmetals. Several metalloids are electrical semiconductors. Elements in a group have similar properties and trends. For example, the alkali metals in Group 1A, alkaline earth metals in Group 2A, oxygen group in Group 6A, halogens in Group 7A, and noble gases in Group 8A all have similar chemical and physical properties within their respective groups. The trends within these groups are similar as well. Alkali metals are soft, metallic solids, and are mostly found in compounds rather than their elemental form. They exhibit typical metallic properties such as luster and conductivity. Alkali metals have low densities, melting points, and ionization energies. By removing an electron, they can achieve the electron configuration of noble gases, which makes it easy for them to form cations with a charge of +1. The chemistry of the alkali metals is characterized by their electronic configuration and various properties such as melting point and ionization energy. Their atomic radius increases down the group, leading to a decrease in melting point. Metal is really exciting, as you can see here. This is lithium in water, sodium in water, and potassium in water. When you add any of those metals to water, they react exothermically, producing a lot of heat. Potassium, for example, can produce white flames, making it highly reactive. The alkaline metals, specifically, react differently with oxygen. Lithium reacts with oxygen to produce lithium oxide, while sodium reacts with oxygen to produce sodium peroxide. The other metals in the same group, when reacting with oxygen, produce superoxide (MO2). So, even though they are in the same group, they react slightly differently in the presence of oxygen. They also undergo flame tests, which are qualitative tests for alkali metals. These metals produce characteristic colors in flames. For example, lithium produces a red color, sodium produces a yellow color, and potassium produces a purple color. This qualitative test helps determine which alkali metal is present. Now, let's talk about the alkaline earth metals. Compared to the alkali metals, the alkaline earth metals have a different electronic configuration. They have the electron configuration of the noble gases. In group 2, the electronic configuration is [noble gas] ns2. They can form cations with a +2 charge by losing those two electrons. This is why their electron configuration is [noble gas] 2+ when they form cations. The alkaline earth metals have different physical properties, such as melting points, densities, atomic radius, and first ionization energy. They have higher densities and melting points compared to the alkali metals, and their ionization energy is low, though not as low as that of the alkali metals. Beryllium does not react with water, and magnesium only reacts with steam. However, the other alkaline earth metals readily react with water. Beryllium and magnesium are relatively more stable compared to the other alkaline earth metals. Moving on to group 6A, the oxygen family. The elements in this group are oxygen, sulfur, selenium, tellurium, and polonium. Oxygen, sulfur, and selenium are nonmetals, tellurium is a metalloid, and polonium is a metal. The first three elements in this group gain two more electrons to achieve a full outer shell, similar to the noble gases. This allows them to have a stable configuration, resembling the next noble gas in the periodic table. Finally, we have group 7A, the halogens. Halogens are typically nonmetals and have highly negative electron affinity. They readily form anions in nature and react with metals to form metal halides. By gaining one more electron, they can achieve the electron configuration of the next noble gas. So, in summary, the alkali metals, alkaline earth metals, oxygen family, and halogens have distinct properties and chemical behaviors. Understanding these groups helps us categorize and study the elements in the periodic table. One key group in the periodic table is Group 8, also known as the noble gases. These gases have unique properties that make them stand out. Firstly, they have very high ionization energy, which means they don't easily lose or gain electrons. Additionally, their electron affinities are positive, making it difficult for them to form stable anions. This is why noble gases are relatively unreactive and nonreactive. In fact, they are often found as monoatomic gases in nature. And with this, we conclude Chapter 7: The Periodic Properties of the Elements.