In 2020, the Spitzer Space Telescope was retired by NASA after 17 years of service. It had observed the universe's infrared activity, providing us with a unique insight into previously hidden areas of space. It carried instruments that could detect wavelengths all the way from just 3.6 μm up to 160 μm in length, and used a mirror almost 1 metre in diameter to focus and reflect light. This mirror was cooled to a chilly 5.5 K - that's -268 °C!
But that's not the reason why it catches our interest. No, we care more about what it was built from. The mirror was made of beryllium, an example of a group 2 element.
Group 2 elements are those in the second column of the periodic table, specifically beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). Group 2 elements are shiny, silvery-white in appearance, and possess relatively low densities. They are highly reactive, although not as reactive as the alkali metals in Group 1. They readily lose two electrons to form +2 cations, resulting in stable electron configurations.
- This article is about group 2 in inorganic chemistry.
- We'll start by providing an overview of the group 2 elements.
- We'll then look at the properties of group 2 elements, including their atomic radius, first ionisation energy, solubility and reactivity.
- After that, we'll consider the uses of group 2.
- Finally, we'll go over how to test for group 2 elements.
Group 2 metal elements
Group 2 is a group of metals in the periodic table. They are also known as the alkaline earth metals.
As mentioned above, group 2 contains six elements:
- Beryllium (Be)
- Magnesium (Mg)
- Calcium (Ca)
- Strontium (Sr)
- Barium (Ba)
- Radium (Ra)
Radium is extremely radioactive, and only occurs as part of the decay chains of heavier elements such as thorium and uranium. Almost all of the naturally-occurring radium in the environment is 226Ra, an isotope with a half-life of 1600 years. However, it's not very common. One kilogram of the Earth's crust contains just 900 picograms of radium - that's 9 x 10-10 grams!
Radium's only current commercial applications are its uses in nuclear medicine, where it can be used to treat certain types of cancers. However, in the early twentieth century, it rose to fame as a source of radiation for radioactive quackery. This is a pseudoscience that improperly promotes radiation as a cure for many illnesses. To this day, you can still find spas that proudly advertise their radium-containing waters as a treatment for all manner of ills and ailments.
In contrast, the group 2 metal calcium is the fifth most common element in the Earth's crust. It has many applications, such as in the production of soaps and cement. However, its most important function is arguably in the body. Calcium is an essential element for many organisms. For example, calcium ions help regulate muscle contraction and nerve function in animals. Our bones act as stores of these ions. A calcium deficiency can lead to osteoporosis. Calcium ions also play a structural role in plants, helping form the cell wall, cell membrane, and middle lamella.
You can find out more about the effect of calcium ions in Sliding Filament Theory.
Group 2 in the periodic table
Take a look at the periodic table below. The column in green shows you one particular group, group 2. As already mentioned, the highlighted elements are beryllium, magnesium, calcium, strontium, barium and radium.
Fig. 1 - Group 2 in the periodic table
Properties of group 2
Group 2 elements are fairly similar. Physically, they are all soft, shiny, silvery-white metals, with relatively low melting and boiling points and densities. Let's take a look at some of their other properties in more detail.
Group 2 structure and bonding
All group 2 elements have two electrons in their outer shell. These electrons are found in an outer s-orbital.
Not sure what we're talking about? Check out Electron Configuration to find out more about different electron orbitals.
When they react, group 2 elements lose their two outer electrons to form cations with a charge of 2+, and an oxidation state of +2. This means that group 2 elements form ionic compounds.
There's one exception to the rule - beryllium. This element actually forms covalent molecules, not ionic compounds. We'll look at why this is so when we move on to group 2's trend in electronegativity.
Group 2 atomic radius
If you've read Periodic Trends, you should be able to predict how the atomic radius of group 2 elements varies as you move down the group. As you can see in the graph below, atomic radius increases moving down the group. This is because each subsequent element has more electrons, with more electron shells.
Fig. 3 - The atomic radius of group 2 elements
We've already seen the electronic structure of magnesium: it has 12 electrons found in three electron shells. The next element in the group, calcium, has 20 electrons found in four electron shells. It, therefore, has a larger atomic radius.
Fig. 4 - Electronic structure and atomic radius of magnesium and calcium
Group 2 melting points
In general, the melting points of group 2 elements decrease as you move down the group. As solids, metals form metallic lattices consisting of positive metal cations surrounded by a sea of negative delocalised electrons, as shown below.
Fig. 5 - Calcium's metallic lattice
This lattice is held together by strong electrostatic attraction between the negative electrons and the nuclei of the positive cations. Remember that atomic radius increases as you move down the group. This means that the nuclei are further away from the delocalised electrons. Therefore, the electrostatic attraction is weaker. So, less energy is needed to overcome it and melt the solid.
Fig. 6 - The melting point of group 2 elements
You'll notice that magnesium's melting point doesn't fit the overall trend. Unfortunately, there's no simple explanation for this. Likewise, the boiling points of group 2 metals don't show a clear trend either. Once again, there's no simple explanation. Yes, we know - deeply annoying!
Need more information on metallic lattices? Metallic Bonding has got you covered!
Group 2 first ionisation energy
We'll now move on to looking at the first ionisation energies of group 2 elements.
First ionisation energy is the energy needed to remove one mole of the most loosely held electrons from one mole of gaseous atoms. Each atom forms a cation with a charge of +1.
First ionisation energy decreases as you go down group 2. Once again, this is due to increasing atomic radius. As you move down the group, the outermost electron is further away from the nucleus. This means that the attraction between the nucleus and the electron is weaker, hence easier to overcome.
Fig. 7 - The first ionization energy of group 2 elements
This topic is covered in much more depth in Trends in Ionisation Energy.
Group 2 electronegativity
Now let's look at electronegativity.
Electronegativity is an atom's ability to attract a bonding pair of electrons.
Once again, there is much more detail in Polarity. But the basic principles of electronegativity apply here too. Electronegativity decreases as you go down the group in the periodic table. As we know, atomic radius increases as you move down the group. This means that any bonded electrons are further from the nucleus, so the attraction between them is weaker.
You might also remember from Polarity that electronegativity is affected by nuclear charge - the number of protons in an atom's nucleus. As you go down the group, nuclear charge increases, so you might think that electronegativity would increase as well.
To explain this, go back to the structures of magnesium and calcium. Magnesium, with an atomic number of 12, has 12 protons in its nucleus. Calcium, on the other hand, has 20. However, magnesium has 10 inner shell electrons that shield the charge of 10 of these protons. In contrast, calcium has 18 inner shell electrons that shield the charge of its protons. In both elements, any bonding pair would therefore only feel the attraction of the two remaining unshielded protons. The effective nuclear charge is the same. But because calcium has a larger atomic radius, it has a lower electronegativity.
Remember that we mentioned that beryllium acts a bit strangely? It forms covalent molecules instead of ionic compounds. This is because it is such a small atom; thus, it has a higher electronegativity than all the other members of the group.
For example, take beryllium chloride and magnesium chloride. Chlorine is much more electronegative than magnesium, and a large difference in electronegativity causes an ionic bond. Chlorine atoms attract magnesium's electrons so strongly that magnesium gives them up completely. Both elements form ions.
Fig. 8 - Magnesium chloride - an ionic compound
On the other hand, beryllium's electronegativity is high enough that it doesn't want to lose its electrons. Instead, it hangs on to them and shares them with chlorine in a covalent bond. This is why beryllium forms covalent molecules instead of ionic compounds.
Fig. 9 - Beryllium chloride - a covalent molecule
Group 2 solubility
Like all metals, group 2 elements are insoluble in water, but their hydroxides and sulphates can dissolve in water to some extent. In particular, hydroxides become more soluble as you go down the group, whilst sulphates become more soluble as you go up the group.
You can find out more in Group 2 Compounds, but here's an overview:
Group 2 reactivity
The final property we'll look at is reactivity. Like most metals, group 2 elements are fairly reactive. Their reactivity increases as you go down the group. As we explored earlier, group 2 elements (apart from beryllium) always react to form ions with a charge of 2+. This requires removing two outer shell electrons - in other words, the processes of first and second ionisation. Ionisation energy decreases as you go down the group, so it is easier to remove these electrons. Therefore, reactivity increases.
We explore some of the characteristic reactions of group 2 metals in Group 2 Reactions.
Testing for group 2
Right at the start of the article, we mentioned how all group 2 elements are pretty similar in appearance. They are all silvery metals. This can make them quite tricky to tell apart. However, one way of distinguishing group 2 metals is by using flame tests. The different metals burn to produce different-coloured flames in a spectacular show of light.
Get a clean metal loop and dip it in acid. Hold the loop in a Bunsen burner flame until there is no colour change. This cleans the loop. Next, dip the loop in a solid sample of your metal and hold it back in the Bunsen burner once again. Observe the colour of the flame produced. With any luck, you'll get the following results:
| Metal | Colour |
| Calcium | Orange-red |
| Strontium | Red |
| Barium | Green |
Note that beryllium and magnesium don't produce a coloured flame. You'll have to rely on other chemical tests to tell them apart.
Uses of group 2 elements
Lastly, let's focus on some of the uses of group 2.
- Calcium is the fifth-most abundant element in the human body and plays a role in bone health, muscle contraction, and neurotransmission.
- Calcium compounds are used in agriculture to raise the pH of soil. They can also be used to remove sulphur from flue gas.
- Barium compounds are used in x-rays, and beryllium alloys are used in mechanical parts.
- Magnesium is the third-most commonly used structural metal, mostly used in lightweight alloys.
Check out Group 2 Compounds for more uses of group 2.
Group 2 - Key takeaways
- Group 2, also known as the alkaline earth metals, is a group of metals in the periodic table.
- Group 2 contains the elements beryllium, magnesium, calcium, strontium, barium, and radium.
- Group 2 elements each have two electrons in their outer shell.
- Atomic radius, reactivity, and solubility of group 2 hydroxides increase as you move down the group.
- Melting point, first ionisation energy, electronegativity, and solubility of group 2 sulphates decrease as you move down the group.
- You can distinguish between some of the group 2 elements using flame tests.
Explanations 
Exams 
Magazine 
