Lesson Video: Periodicity | Nagwa Lesson Video: Periodicity | Nagwa

Lesson Video: Periodicity Chemistry • Second Year of Secondary School

In this video, we will learn how to describe the trends in properties across periods of the periodic table.

14:45

Video Transcript

In this video, we will learn how to describe the trends and properties across periods of the periodic table.

The periodic table is an amazing display of information. Simply by listing the elements in order of their number of protons and starting new rows at the right time, we can arrange all of the elements into a visual representation that reveals important information and trends. In chemistry, periodicity is the idea that there are trends and patterns that we can find in the periodic table. For the simplest trends, we will see whether a certain characteristic increases or decreases as we move in two directions in the periodic table. We can look at the trend down the group, and we can also look at the trend across the period. In this video, we will take a look at the periodicity of four different characteristics: atomic radius, ionic radius, melting point, and electrical conductivity.

First, let’s take a look at atomic radius. If we consider the structure of an atom with a central nucleus and an outer ring of electrons, the atomic radius is the distance between the nucleus of the atom and its outer electrons, the radius of the atom. However, it’s hard to directly locate the outermost electron for this measurement, so chemists measure the atomic radius a different way. They take two atoms bonded to one another and measure the distance between two nuclei to find a distance equal to two atomic radii. For example, the distance between adjacent nuclei in a chlorine gas molecule is 202 picometers. Since we consider the two molecules to be adjacent with no space between them, the atomic radius of chlorine is half of that value or 101 picometers.

So, what are the periodic trends for atomic radius? The first trend is that atomic radius increases as we go down a group. To investigate this trend, let’s take a look at the electron shells of two elements in the same group. Sodium is lower than lithium in group one. As a result, it has more electrons. And those electrons fill more electron shells. The end result is that sodium has a larger atomic radius than lithium. And in fact, in any group on the periodic table, as we move down the group, we will increase in atomic radius. The trend in atomic radius in the other direction is that as we move from left to right within a period, the atomic radius decreases. To investigate this trend, let’s take a look at two elements in the same period.

Nitrogen is to the right of carbon in period two. Therefore, it has more protons. More protons means more positive charge, which means a stronger pull on the outer electrons of the atom. The extra inward pull on the electrons reduces the atomic radius. So, nitrogen has a smaller atomic radius than carbon. And as we move to the right in a period on the periodic table, the atomic radius gets smaller and smaller. The second pattern sometimes seems counterintuitive. At first glance, you might think that adding protons and electrons could increase the atomic radius.

As a metaphor to help us understand the second trend, we can think of the atom as a concert with a popular proton boy band on stage with an audience of adoring electron fans. The concert has a certain radius, with fans standing in neat rows. Adding more fans to an existing row doesn’t change the radius of the concert. But if another boy band member emerges on stage, the adoring electron fans might move in closer to get a better look, taking up less space as a result. It’s only when we add fans to a new row or add electrons to a new electron shell that the radius of the concert increases. With this metaphor in mind, we can better understand why adding more protons and electrons increases the attractive forces between them and decreases the size of the atom.

For the periodic trend, we can look at two ions from the same period of the periodic table, such as an oxygen ion and a fluoride ion. Whether these particles are ions or atoms, the one with more protons will have a stronger pull on the electrons and thus a smaller atomic or ionic radius. We can extend this pattern a bit further if we look at a group of ions or atoms that are isoelectronic. Two or more atoms or ions are said to be isoelectronic if they have the same number and arrangement of electrons. This group of atoms and ions here is isoelectronic. They all have 10 electrons, completely filling the first two electron shells. The main difference between these ions and atoms is the number of protons.

Nitrogen on the left has seven protons and the protons increase by one up to magnesium on the right, which has 12 protons. The same rule from earlier still applies. More protons means more pull on the outer electrons, which means a smaller atomic radius. So, as we look at the isoelectronic atoms and ions with more and more protons, the atomic radius decreases. So the second pattern holds for ions within a period, as well as isoelectronic ions across several periods.

Another characteristic we can investigate is melting point. Melting a substance requires adding enough energy to overcome the attractive forces holding the solid structure together. So, a high melting point is an indication that the substance contains strong attractive forces holding its particles together. This relationship is apparent when we compare metals and nonmetals. Metals are held together by strong metallic bonds that take a lot of energy to break. As a result, in general, metals have high melting points. Nonmetals, on the other hand, are held together by the much weaker London dispersion forces. As a result, in general, nonmetals have lower melting points than metals.

This pattern matches what we observe in real life. At room temperature, the vast majority of metals are solid, meaning that their melting point is higher than room temperature. The vast majority of nonmetals, on the other hand, including oxygen gas and nitrogen gas, are gaseous at room temperature. For these substances, their melting points as well as their boiling points are below room temperature. It’s worth noting that this is a general trend. There are some specific exceptional values that may deviate from this pattern.

Let’s investigate melting point more precisely by taking a closer look at one period of the periodic table. The third period of the periodic table contains the elements sodium, magnesium, aluminum, silicon, phosphorus, sulfur, chlorine, and argon. Starting with the elements to the left, we have sodium with a melting point of 98 degrees Celsius, magnesium with a melting point of 639 degrees Celsius, and aluminum with a melting point of 660 degrees Celsius. These three elements are metals with relatively high melting points.

As we move to the right among these three elements, we increase the number of free-floating electrons in the metallic bond. Sodium has one free-floating valence electron, magnesium has two free-floating valence electrons, and aluminum has three free-floating valence electrons. These negatively charged particles add extra attraction and therefore strength to the metallic bonds that make up the metal substances. With stronger attractive forces holding the substance together, the melting point of that substance is higher.

Silicon has the highest melting point of any element in period three at 1410 degrees Celsius. This high melting point of silicon is due to its structure, specifically its giant covalent structure. A giant covalent structure is held together by many covalent bonds. Each silicon atom is covalently bonded to four other silicon atoms in its immediate vicinity. The number and strength of the bonds in the structure means that it requires lots of energy to break apart the solid structure of the substance. As a result, the melting point of silicon is very high.

Looking at the remaining elements to the right, no clear pattern emerges. In general, we just need to recognize that these are nonmetals and therefore they have lower melting points. These melting points are lower because the particles within the substance are held together by the weaker London dispersion forces. Although it is worth noting that the nonmetal sulfur does have a higher melting point than the metal sodium. As a reminder, these are just general trends. Other characteristics beyond metallic character can influence the melting point.

Another characteristic relevant to metals and nonmetals is electrical conductivity. Electrical conductivity is the ability of a substance to carry an electric current since electricity is the flow of charged particles. To know how conductive a substance is, we need to ask how free are the electrons to move. On the left side of the periodic table, we have metals. Metals are held together by metallic bonds. Metallic bonds have delocalized electrons. In other words, some of the electrons float freely among the atoms of the substance. As a result, metals have high electrical conductivity.

On the other side of the periodic table, we have nonmetals. Nonmetals tend to form covalent bonds, although some nonmetals exist as monatomic elements, in other words, as lone atoms without any bonds. While the electrons of these bonds do move around, they’re more confined to the space between atoms rather than flowing freely about the substance. As a result, nonmetals have low electrical conductivity.

In between the metals and nonmetals on the periodic table are the semimetals. Semimetals sometimes behave like metals and sometimes behave like nonmetals. So, semimetals have medium electrical conductivity. By visually separating metals from nonmetals, the periodic table helps us see this trend in electrical conductivity. Now that we’ve learned about periodicity, let’s do a practice problem to review.

Which of the following elements has an atomic radius larger than that of aluminum? (A) Chlorine, (B) sodium, (C) sulfur, or (D) silicon.

This question is asking us to compare the atomic radii of different elements. The atomic radius of an element is the distance from its nucleus to its outermost electrons. Scientists can measure the atomic radius by taking half the distance between the two nuclei of adjacent atoms bonded together. To answer this question, we need to take a look at the periodic table, specifically at period three, which contains all five of the elements mentioned in the question.

In order to answer this question, we need to know how atomic radius changes across a period. Once we know this trend, we can find which of the answer choices has the largest atomic radius. Let’s compare aluminum and silicon to visualize this change. Since we don’t know what size they are yet, we’ll start by drawing them as the same size. The first difference we can notice is that when we move to the right in a period, the number of protons increases. Silicon has one more proton than aluminum. With more protons in the nucleus, there’s a stronger positive charge there and thus a stronger inward pull on the electrons. As a result, there’s also a smaller atomic radius.

So, the trend is clear. If more protons means a smaller atomic radius, then the atomic radius decreases to the right and increases to the left. If we wanna find the element with an atomic radius larger than that of aluminum, we want to find the element to the left of aluminum on the periodic table. In this case, that’s sodium. Sodium has fewer electrons than aluminum and therefore a weaker inward pull on its electrons, which results in a bigger atomic radius. Which of the following elements has an atomic radius larger than that of aluminum? That’s answer (B) sodium.

Now that we’ve learned about periodicity, let’s review the key points of the video. The word periodicity refers to the trends and patterns found in the periodic table. One trend is that atomic radius increases down a group and decreases across a period. These trends are due to the layering of electron shells and the strength of the attractions between the nucleus and the electrons, respectively.

Melting point is an indicator of the strength of the attractive forces that hold the substance together. In general, metals held together by strong metallic bonds have higher melting points than nonmetals, which are held together by the weaker London dispersion forces. Another characteristic, electrical conductivity, depends on the free movement of electrons. Metals are more conductive than semimetals, which are more conductive than nonmetals.

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