Lesson Video: Physical Properties of Transition Metals Chemistry • Third Year of Secondary School

Video Transcript

In this video, we will learn about the physical properties of the transition metals, specifically the first row of the transition metals. Let’s start by briefly reminding ourselves what a transition metal is and where these metals are located on the periodic table.

Transition metals are those elements whose atoms have an incomplete d subshell or whose cations have an incomplete d subshell. They are located in the d-block of the periodic table, and the inner transition elements are located in the f-block. We will only study the first row of the main transition metals. This row is called the first transition series, and it is in period four. The transition metals in the series are scandium, titanium, vanadium, chromium, manganese, iron, cobalt, nickel, and copper. Note that zinc in group 12 is not technically considered to be a transition metal because its d subshell is complete and full with electrons.

Now you might be familiar with some notable physical properties of some of these elements. For example, titanium is low in density but very strong, polished chromium is highly lustrous or shiny, and copper is an excellent electrical conductor. Before we have a look at some specific physical properties of the first transition series of elements, let’s have a look at their atomic masses.

The atomic masses for these elements are listed and given in atomic mass units. Going from left to right across the period, we find a general trend: as the atomic number or number of protons increases, so the atomic mass increases too. However, there is a slight dip in the trend for nickel. Let’s see how the atomic mass influences the atomic radius.

Now the atomic radius is the distance from the center of the nucleus to the outermost shell or the outermost electron-containing shell. If this picture is an atom, this is the nucleus, and this is the outer electron-containing shell, then the pink arrow represents the atomic radius. Atomic radii can be measured in different ways. Depending on how they are measured will give slightly different values. And they can be measured using different units, for example, in picometers, which is one times 10 to the negative 12 meters, or in angstroms, which is one times 10 to the negative 10 meters. The values shown here are in angstroms.

Moving from left to right across the series, we notice a general trend: as the atomic number or number of protons increases, so the atomic radius decreases with just a slight increase at the end for copper. Why is there this general trend? Well, the more protons there are in the nucleus, the higher the nuclear charge, in other words, an increased attractive force on the outer electrons. And so, the outer electrons are held closer or pulled closer to the nucleus, resulting in a smaller atomic radius. If we plot these values on a graph of atomic radius versus increasing atomic number, the curve would look something like this. It is interesting to note that the bottom of the curve between chromium and copper is fairly flat. This is because there is increasing repulsion between the 3d electrons and this counteracts the attractive force of the nucleus. We say there is a lower effective nuclear charge.

It is interesting to note that potassium and calcium, which come just before scandium in period four in the periodic table and which are shown here on the graph too, have significantly larger atomic radii. We won’t discuss why this is so in this video. But we can note that potassium and calcium are very different species to these transition metals because potassium and calcium do not have d orbitals. Now we know about the atomic radius trend. But what about density?

The densities of these period four elements are shown here. They are given in grams per centimeter cubed, and the values are roughly plotted here on a graph to show us the general trend visually. We can see that going from left to right across the series as the atomic number or number of protons increases, so the density increases. What about the s-block elements potassium and calcium? How do their densities compare with these transition metals?

Potassium’s density value, which is about 0.86 grams per cm cubed, and calcium’s value, which is about 1.55 grams per cm cubed, are notably lower than the values for all the other transition metals. We know that the outer electronic configuration of these transition elements is 4s2 3dx, where x is a value of one to a value of nine depending on the element. Potassium and calcium only have the 4s electrons available for metallic bonding. But the first transition series has both the 4s electrons as well as the 3d electrons available for metallic bonding, and the strength of the metallic bonds influences the density.

The property up next is melting point. Melting point does not have a definitive trend for the first transition. series. From scandium’s melting point of 1397 degrees Celsius, the values go up, up, up, down, up, down, up, down. I won’t plot these values on a graph for you. The graph kind of looks like a roller coaster. Boiling point values for the series also don’t follow a recognizable trend. However, let’s have a look at the boiling points of the s-block elements just before these metals.

Potassium’s value is 63.5 degrees Celsius and calcium 842 degrees Celsius. So, comparing these s-block element melting point values with these, transition metal values, we can make a conclusion. We can see that the fourth period transition metal melting points are much larger than the fourth period s-block element melting points. Why is this?

As before, the transition metals have both the 4s electrons and the 3d electrons available for metallic bonding. And the s-block elements only have the 4s electrons available for metallic bonds. So, the transition metals have stronger metallic bonds, and the s-block metals have weaker metallic bonds. Stronger metallic bonds result in higher melting points, and weaker metallic bonds result in lower melting points. What is interesting to note is that calcium’s melting point is higher than zinc’s melting point, even though zinc has 3d electrons available for metallic bonding. We will not go into an explanation of this here, but I will remind you that zinc is technically not a transition metal.

Let’s move on to the magnetic properties of these elements. Now some atoms, ions, or molecules are attracted to a magnetic field, and some are not. Those particles that are attracted to a magnetic field are called paramagnetic particles. Those that aren’t attracted are called diamagnetic. Diamagnetic particles are actually slightly repelled by a magnetic field. Paramagnetic particles have unpaired electrons in the orbitals. These unpaired electrons create a magnetic field due to their spin. The more unpaired electrons there are, the larger the magnetic moment of a particle.

Diamagnetic atoms or particles have no unpaired electrons; all their electrons are paired. Electrons pair with opposite spins. So, in a pair of electrons, each electron will negate the effect of the spin of the other electron. In other words, a pair of electrons has no net electron magnetic dipole moment, and so no magnetic field is created by a pair of electrons with opposite spins. This is why particles whose electrons are all paired are not attracted to a magnetic field but are rather just slightly repelled.

Let’s look at an example. This is scandium’s electronic configuration. The two 4s electrons are paired with each other, but the lone or single 3d electron has no partner. It is unpaired. So, an atom of scandium is paramagnetic. Here’s another example. This is the electronic configuration of an atom of iron. Again, the two 4s electrons are paired with each other. But what about the 3d electrons? If we put the six 3d electrons into the five orbitals of the 3d subshell, we see that two of the electrons are paired and four are unpaired. So, an atom of iron is paramagnetic and to a larger degree than an atom of scandium because an atom of iron has more unpaired electrons.

Now, let’s change this atom of iron to an ion of iron. Let’s make it Fe3+. The electron configuration of Fe3+ is shown here. Three electrons have been removed from the Fe atom to create the Fe3+ ion, two electrons from the 4s subshell and one from the 3d subshell. We can put the electrons in the 3d subshell into separate orbitals as shown. And we can determine that Fe3+ is also paramagnetic because it contains unpaired electrons, in this case, five unpaired electrons. What about zinc? If we write the electronic configuration of zinc, we will see that there are 10 3d electrons, in other words, five pairs of electrons. Zinc has no unpaired electrons, and it is diamagnetic.

We’re going to look at two more properties of these transition metals. Next up is catalytic activity. Many transition metals are very useful in industry due to the excellent catalytic activity. For example, iron is used in the Haber–Bosch process. This is the production of ammonia. Vanadium is used in the contact process in the form of V2O5 for the production of sulfuric acid. Nickel is used in a variety of hydrogenation reactions. An example of a hydrogenation using Raney nickel is the hydrogenation of benzine into cyclohexane transition. Metal catalysts and their catalytic activity are covered in more depth in another video.

We will look at the last property now, and that is the color of transition metal complex compounds. When white light, which is composed of the seven colors of the rainbow, passes through an aqueous solution of a complex transition metal compound, some wavelengths are absorbed, and some will pass through the solution. For example, this complex transition metal compound when in aqueous solution absorbs the red and orange wavelengths of white light. However, yellow, green, blue, indigo, and violet pass through without being absorbed. The mixture of these emerging wavelengths appear pale blue-green or cyan to the human eye.

This handy color wheel shows us that if red and orange wavelengths are absorbed, then the opposite or complementary colors on the wheel — in this case, blue and green — will be the appearance of the solution to the human eye. Cyan is a pale blue-green color. Don’t get confused. Violet, indigo, and yellow wavelengths will be emerging too. This color wheel merely shows us what the solution will look like to the human eye. The particular wavelengths absorbed depend on the electrons in partially filled d orbitals. Further explanation of this is covered in another video.

Now, it’s time to wrap up by summarizing what we have learnt. In the first transition metal series, the atomic mass increases as we go from left to right, with the exception being nickel. The atomic radius decreases as the nuclear charge increases when we go towards the right. Density values increase the further we go to the right. The melting point values have no trend but are higher than the period four s-block elements potassium and calcium. We said this is because of the increased metallic bonding in these transition metals due to the presence of 3d electrons.

We saw that all atoms of the first transition metal series are paramagnetic due to the presence of unpaired electrons. The exception is zinc, whose electrons are all paired. However, zinc is not technically a transition metal according to the definition. We looked at just three examples of transition metals used as catalysts. Lastly, we learnt that the color of complex transition metal compounds in aqueous solution depends on the 3d electrons and the corresponding wavelengths absorbed. We learnt that the human eye detects the complementary colors, those that are not absorbed by the solution.