Lesson Video: Group 15 Chemistry

In this video, we will learn the Group 15 elements: their elemental properties, the reasons for those properties, and their economic importance.


Video Transcript

In this video, we will learn about the group 15 elements, their elemental properties, the reasons for those properties, and their economic importance. Group 15, otherwise known as group five, is a collection of elements on the periodic table. They’re sometimes known as the pnictogens. It’s a collection of elements whose atoms have a very similar electronic configuration. There are currently six elements in total in group 15. At the top is nitrogen with atomic number seven. Then there’s phosphorus with atomic number 15, arsenic with atomic number 33, antimony with atomic number 51, bismuth with atomic number 83, and finally moscovium with atomic number 115.

Phosphorus and nitrogen are the only nonmetals in the group. But nitrogen is the only one that’s commonly found as a gas at room temperature in the form of the diatomic N2 molecule, while phosphorus forms a variety of molecular and polymeric solids. Meanwhile, arsenic and antimony have properties in between those of metals and nonmetals, so they’re generally classified as metalloids. These typically form solid lattices that are insulators, conductors, or semiconductors, although arsenic also has a molecular form. And bismuth is classed as a metal, although it’s classed as a poor metal since it has a very low conductivity. Nonetheless, it exhibits metallic bonding. Lastly, moscovium is a synthetic element. Not enough has been made for us to assess its chemistry and its physical properties.

As we descend group 15, each element is progressively less abundant on Earth than the previous one. Nitrogen is the most readily available as N2 molecules make up 80 percent of the molecules in the atmosphere. The remaining elements are found in compounds in the Earth’s crust. About 1000 milligrams out of every kilogram of the Earth’s crust is made up of phosphorus. We can also express this is about 0.1 percent by mass. By mass, arsenic is about 500 times less abundant than phosphorus, making up only about two milligrams per kilogram of the Earth’s crust.

Antimony is 10 times less abundant still. And relatively speaking, bismuth is relatively rare, making up only 0.01 milligrams per kilogram. We can look at aluminum to put these numbers in context. Aluminum has an abundance in the Earth’s crust, about 80000 milligrams per kilogram. Of course, we don’t have a natural abundance for moscovium since it’s entirely synthetic. For the remainder of the video we’ll only be looking at the top five members of group 15.

Before we move on, we’re going to have a look at what natural forms these elements are found in. The most common form for nitrogen, as we’ve already said, is N2 gas, but we can also find it as nitrates in soil and in other compounds. Nitrogen is the only element in group 15 that’s readily found in nature in its elemental form. The rest are found as compounds. Phosphorus is generally found bound up in phosphates such as calcium phosphate and minerals such as apatite, while arsenic, antimony, and bismuth are generally found bonded to sulfur in sulfides.

Another interesting feature of these elements is the range in forms they exhibit when they’re turned into a gas. Of course, we know nitrogen forms N2 molecules, but phosphorus, arsenic, and antimony form tetrahydral four-atom molecules. And uniquely for a metal, bismuth will form a Bi2 molecule in the gas phase. The next thing we’re going to look at is the electronic configuration of atoms of these elements.

This is where the magic happens. As we descend the group, atoms of the elements have increasing nuclear charge and increasing numbers of electrons in the electron cloud. There’ll be seven electrons in the electron configuration for nitrogen, 15 for phosphorus, 33 for arsenic, 51 for antimony, and 83 for bismuth. The highest energy subshell we’ll need to populate is the 6p subshell in a bismuth atom. So that’s the extent I’ve taken the Aufbau diagram. If we use the Aufbau principle and populate the lowest-energy orbitals first, we’ll be able to figure out the electron configurations for these atoms.

The ground-state electron configuration of a nitrogen atom is 1s2 2s2 2p3. That’s seven electrons in total. The 2p subshell here is half full. The extra eight electrons that our phosphorus atom has over a nitrogen atom go into filling the 2p sub shell, the 3s subshell, and half filling the 3p subshell. When we complete the electronic configuration of arsenic, we need to be careful to add electrons to the 4s subshell before we do the 3d subshell. After the 4s subshell is full, we can add 10 electrons to the 3d subshell and half fill the 4p subshell.

The pattern continues with antimony, with the last three electrons in the electronic configuration of an antimony atom going into the 5p subshell. And the electronic configuration of a bismuth atom is quite long, but the last three electrons do end up in a p-type subshell. Before we do a thorough comparison, let’s condense these configurations. We can condense the core electron configurations by comparing them to noble gases. This way we can focus on the outer electrons. Atoms of nitrogen or phosphorus clearly only have five electrons in their highest n electron shell. We also see the same for atoms of arsenic, antimony, and bismuth.

For elements in group 15, we typically see three or five electrons per atom being involved in bonding. That’s either three electrons being used from the p-subshell or two from the s and three from the p. We can see evidence of this when we look at their common oxidation states. As we descend group 15, we see less variation in the oxidation states. Nitrogen can exhibit oxidation states anywhere between negative three and positive five. We tend to see negative oxidation states when nitrogen is bonded to hydrogen, in ammonia, hydrazine, and hydroxylamine, while the positive oxidation states tend to crop up in compounds containing nitrogen-oxygen bonds from plus one in nitrous oxide to plus five in nitric acid.

Now phosphorus, phosphorus can exhibit oxidation states of negative one or negative two. But the most common oxidation states are negative three, zero, positive three, and positive five. We can see the negative three oxidation state in calcium phosphite, where we see the P3- ion. And we typically see the positive oxidation states when phosphorus is bonded to halides or oxygen. Arsenic exhibits similar behavior showing oxidation states of negative three, zero, positive three, or positive five, while antimony and bismuth in compounds tend to only exhibit the positive three or positive five oxidation states.

Some of the group 15 elements exhibit different pure forms called allotropes. We’ll leave out nitrogen and bismuth since these elements only have one common allotrope. The other allotropes require extreme conditions in order to be stable. The common allotropes of the other three elements can be achieved with very reasonable conditions. There are four well-recognized allotropes of phosphorus: white, red, violet, and black, although black is much harder to obtain. White phosphorus is composed of tetrahydral P4 molecules. White phosphorus is so reactive it’s pyrophoric. It’ll react with oxygen in air and self-ignite. It’s also highly toxic and volatile and relatively unstable, transforming into red phosphorus in response to light and heat.

Red phosphorus is polymeric made up of strings of phosphorus atoms. The conversion of white phosphorus to red phosphorus can be accelerated by heating white phosphorus to 250 degrees Celsius. Violet phosphorus is simply the crystalline form of the red phosphorus polymer. We can convert red phosphorus to violet phosphorus by heating it at or above 550 degrees Celsius. And finally, black phosphorus resembles graphite in its structure. Like graphite, black phosphorus is conductive. Black phosphorus is less readily available than the other allotropes of phosphorus since it requires catalysts or incredibly high pressures to produce from white phosphorus.

Next, let’s look at arsenic. The three most common allotropes of arsenic are gray, yellow, and black. Gray arsenic has a layered structure, and it’s quite brittle. It’s the most stable and the most common of the three. Gray arsenic is a semimetal, and it can be turned into a semiconductor. Yellow arsenic is soft and waxy and is composed of As4 molecules. Like white phosphorous, it’s unstable, volatile, and quite toxic. We can make yellow arsenic by condensing arsenic vapor. And lastly, black arsenic is glassy, brittle, and a poor conductor.

Next, let’s look at antimony. The most common forms of antimony are metallic, black, and yellow. But there’s also an exotic explosive variety. The metallic variety is silvery white and quite brittle. The explosive variety is produced by the electrolysis of antimony trichloride. When scratched, white clouds of metallic antimony are formed, and in a pestle and mortar, it explodes. Black antimony is glassy and brittle, and yellow antimony is only stable at low temperatures below negative 90 degrees Celsius. Above this temperature, it will turn into black antimony. Now we’re going to move on to some of the compounds of the group 15 elements.

We’ll start with the oxides. As we descend the group, the oxides of these elements become more basic. Nitrogen at the top of the group forms the most acidic oxides. All these oxides from nitrous oxide to dinitrogen pentoxide are acidic to some degree. Many of these oxides can be produced by direct reaction of nitrogen with oxygen gas. For example, nitrogen dioxide can be produced by reacting one lot of nitrogen with two lots of oxygen. Nitrogen dioxide can react with water to produce a mixture of nitrous and nitric acids.

Next, let’s look at phosphorus. The common oxides of phosphorus are phosphorus three oxide and phosphorus five oxide. These are both acidic, forming phosphorous acid and phosphoric acid, respectively. White phosphorous reacts with oxygen in the air to form phosphorus five oxide in the form of a white powder. The single P4O10 unit reacts with six water molecules to produce four molecules of phosphoric acid.

Next up, arsenic. Arsenic tends to form arsenic three oxide over arsenic five oxide. But both are acidic. Arsenic three oxide can be produced by burning arsenic three sulfide. Now, antimony. Like arsenic, antimony tends to form the antimony three oxide over the antimony five oxide. These oxides are amphoteric, being basic in some circumstances and acidic in others. And finally, we have bismuth. Bismuth reacts with oxygen to form the bismuth three oxide. Forming the bismuth five oxide is much more difficult. These oxides are basic. Bismuth three oxide will react with water to form bismuth three hydroxide.

Next, we’re going to look at the hydrides. As we descend group 15, the hydrides become less polar. As a consequence, they become less soluble in water. Technically, ammonia is not a hydride. Nitrogen has a negative three oxidation state and hydrogen has a positive one oxidation state. We’ll leave that technicality aside for now. We can predict the oxidation state of the group 15 element using their electronegativities relative to hydrogen. Nitrogen is three, substantially above hydrogens at 2.2. The electronegativities of phosphorus and arsenic are almost exactly the same as hydrogens, while the electronegativities of antimony and bismuth are clearly lower than that of hydrogen.

The lone pair on the nitrogen in ammonia provides an area of concentrated negative charge. In the nitrogen-hydrogen bonds, electrons end up closer to the more electronegative nitrogen atom, further increasing the size of the overall dipole. It’s the large dipole of ammonia that allows it to be highly water soluble. The phosphorus in phosphine, PH3, also has a lone pair, providing a little bit of a dipole. But phosphorus and hydrogen have about the same electronegativity, so the phosphorus-hydrogen bonds are nonpolar and don’t contribute to the overall dipole.

Like phosphine, arsine has a lone pair, but it’s more diffuse. So the overall dipole is smaller. With antimony, for the first time in group 15, we have an element whose electronegativity is lower than that of hydrogens. Hydrogen, therefore, adopts the negative one oxidation state. The hydride of antimony is known as stibine. In stibine, the polarity of the antimony hydrogen bonds are competing against the dipole produced by the lone pair. Overall, stibine has a very low dipole. And the pattern continues with bismuthine.

Now, we’re going to look at the economic importance of each of these elements. The main industrial use for nitrogen is in the production of ammonia, a vital ingredient in the fertilizer industry. But it’s also used to inflate car tires, keep potato chips fresh, and preserve tissue samples and do air-sensitive chemistry. Phosphorus, like nitrogen, is also a vital part of the fertilizer industry. It’s also a key component of match heads and match boxes, in alloys such as phosphorus bronze, and it’s found in some firework ingredients.

Arsenic and many of its compounds are notoriously toxic. Nonetheless, arsenic has a few uses, for example, in wood preservatives, killing fungi, bacteria, and viruses. And it’s even used in medicines. For instance, arsenic three oxide is used to treat leukemia. Antimony can be found in various alloys, for instance, in combination with lead in the alloys in car batteries. And it’s also found in some semiconductors. Bismuth is also found in various alloys, for instance, combined with lead and cadmium to produce low melting point alloys in fuses.

Now it’s time to finish up with the key points. The group 15 elements, otherwise known as group five elements, are nitrogen, phosphorus, arsenic, antimony, bismuth, and moscovium. Their atomic numbers are seven, 15, 33, 51, 83, and 115, respectively. The atoms of these elements all have electronic configurations that finish with an s2 p3 configuration. We only need to explore the electronic configurations of atoms of elements up to bismuth. Moscovium is a synthetic element, and its chemistry has not been thoroughly investigated. In compounds, we see oxidation states for nitrogen between negative three and positive five. As we descend the group, the range of oxidation states decreases and compounds with positive five oxidation state group 15 elements become more unstable.

Phosphorus allotropes include white, red, violet, and black. You can get gray, yellow, or black arsenic. Antimony’s most common allotropes are metallic black and yellow, while nitrogen and bismuth have only one common allotrope. For nitrogen, it’s N2 gas, and for bismuth, it’s a metallic form. The oxides of nitrogen and phosphorus are all acidic. And those of arsenic are more weakly acidic. The oxides of antimony are basic in some circumstances and acidic in others. So they’re amphoteric, while oxides of bismuth are basic. And the compounds of group 15 elements with hydrogen of the form XH3 become less polar and less soluble as you go down the group.

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