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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 is a group of metals in the periodic table. They are also known as the alkaline earth metals.
Take a look at the periodic table below. The column in pink shows you one particular group, group 2.
Group 2 in the periodic table. Olive, StudySmarter Originals
Group 2 contains six elements:
These elements all appear in nature, although barium is never found on its own due to its high reactivity. Likewise, 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 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 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.
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.
The electron configuration of magnesium. Created using images from commons.wikimedia.org
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.
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.
The atomic radius of group 2 elements. Anna Brewer, StudySmarter Originals
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.
Magnesium and calcium. Created using images from commons.wikimedia.org
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.
Calcium's metallic lattice. Anna Brewer, StudySmarter Originals
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.
The melting points of group 2 elements. Anna Brewer, StudySmarter Originals
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!
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.
Can you guess the trend?
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.
The first ionisation energy of group 2 elements. Anna Brewer, StudySmarter Originals
This topic is covered in much more depth in Trends in Ionisation Energy.
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.
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.
Beryllium chloride. Anna Brewer, StudySmarter Originals
Like all metals, group 2 elements are insoluble in water. However, their compounds have varying solubilities. (You can find out more in Group 2 Compounds.)
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.
Now that we've learnt about the properties of group 2, we can take a look at some of their reactions.
All reactions involving group 2 are redox reactions.
Redox reactions are reactions involving both oxidation and reduction. Oxidation is the loss of electrons, whilst reduction is the gain of electrons. (See Redox for more details.)
The group 2 element loses two electrons to form a cation with a charge of 2+, and an oxidation state of +2. It is oxidised. On the other hand, the other elements involved gain electrons. They are reduced. Don't worry - we'll practice showing this below.
Let's now explore the reactions of group 2 with water, oxygen, chlorine, and acids.
We're ignoring beryllium in all of the following examples - it doesn't form ions, as we saw earlier. What a tricky exception to the rule!
Have you ever wondered why group 2 are called the alkaline metals? It's because they react with water to produce alkaline metal hydroxides and hydrogen gas. This is a great example of the trend in reactivity of group 2 elements. Here, M represents the group 2 metal:
Although we've shown the metal hydroxide as aqueous, this might not always be the case. The solubility of group 2 metal hydroxides actually varies, increasing as you move down the group. In fact, magnesium hydroxide is essentially insoluble, whilst calcium hydroxide is only sparingly soluble. Strontium and barium hydroxide are both extremely soluble in water.
As with all reactions involving group 2 elements (apart from that sneaky beryllium), the reaction between a group 2 metal and water is a redox reaction. The group 2 metal is oxidised and loses electrons, whilst the hydrogen in water is reduced and gains an electron. The oxygen doesn't take part in the redox reaction - it is just a spectator species. We can show this using oxidation states.
The oxidation states in magnesium hydroxide. Anna Brewer, StudySmarter Originals
Head over to Redox for more.
The reaction between a group 2 metal and water gets more vigorous as you go down the group. For example, magnesium reacts extremely slowly with water, producing just a few bubbles. You'd be hard-pressed to see any reaction at all. However, calcium fizzes gently when you add it to water, whilst strontium and barium bubble violently. Overall, their reactivity increases as you move down the group. Once again, this is because their ionisation energies decrease as you go down the group. Lower ionisation energies mean that it is easier to lose electrons, so reactivity increases.
We'll now focus on the reaction between group 2 elements and oxygen. Burning a group 2 element in oxygen produces a metal oxide. Here's the equation:
The same reaction happens if you leave a group 2 metal exposed to air. A metal oxide rapidly builds up on the surface of the metal, producing a thin coating that stops any further reaction.
Strontium and barium also burn in oxygen to produce a peroxide, . However, beryllium, magnesium, and calcium don't. This has to do with the smaller size of these metal ions. Since they are smaller, they have a much higher charge density, attracting the oxygen molecules differently.
Metal peroxides are formed when the peroxide ion, O22-, reacts with a metal ion. The peroxide ion consists of two oxygen ions, each with a negative charge or 1-, covalently bonded together.
All elements in group 2 form ions with a charge of 2+. However, beryllium, magnesium, and calcium form smaller ions than strontium and barium. This means that they have a higher charge density. The charge density is so large that it is able to pull some of the electrons in the peroxide ion over to one of the oxygen atoms involved, breaking the covalent bond between them. This results in a neutral oxygen atom and an O2- ion, which bonds with the metal ion.
However, strontium and barium are larger ions. This means that they have a lower charge density. Their charge density isn't great enough to disrupt the covalent bonding in the peroxide ion. The peroxide ion remains intact when it bonds to the metal, so strontium and barium are able to form peroxides.
Magnesium also reacts with steam to produce the same product, magnesium oxide. The reaction initially forms magnesium hydroxide, but this splits up upon heating to produce the oxide and hydrogen gas. Overall:
Group 2 metals react with chlorine to produce metal chlorides. In fact, they react in a similar way with all halides. Once again, reactivity increases as you go down the group. Here's the general equation:
Finally, we'll turn our attention to the reactions of group 2 with acids. These vary depending on the type of acid used.
This is the simplest case - the group 2 element reacts with hydrochloric acid to give a metal chloride and hydrogen gas. For example, mixing magnesium with hydrochloric acid gives magnesium chloride and hydrogen:
Magnesium chloride is the primary salt found in the Dead Sea. With a salinity of over 30 percent, it is one of the world's saltiest bodies of water.
Nitric acid makes things a little more complicated. Group 2 metals can also reduce the nitrogen atoms found within nitric acid, forming all manner of nitrous oxides. However, the main reaction produces a metal nitrate and hydrogen gas.
Group 2 elements react with sulphuric acid to produce metal sulphates and hydrogen gas. However, in reality, strontium and barium don't tend to react much. This is because the sulphate forms an insoluble layer on the surface of the metal, preventing any further reaction. Likewise, calcium sulphate is only sparingly soluble, so calcium doesn't react much with sulphuric acid either. Nevertheless, here's the general reaction between a group 2 metal and sulphuric acid:
We look at the solubility of metal sulphates in more detail in Group 2 Compounds.
One further reaction involving group 2 metals uses magnesium to extract titanium. Discovered in Cornwall in 1791 by William Gregor, titanium is an extremely useful transition metal. It has a low density but high strength, and is resistant to both seawater and chlorine. You'll find it in the aerospace industry as well as in mobile phones and orthopaedic implants.
Titanium is found deep within the Earth's crust as titanium oxide. This is first reacted with chlorine to produce titanium tetrachloride, . Magnesium is then added. This reduces the titanium within titanium tetrachloride, forming magnesium chloride and titanium metal:
Here's a handy table that summarises the reactions of group 2 for you.
Element | Reactant | Product | Equation |
All group 2 elements | Cold water | Metal hydroxide and hydrogen | |
Magnesium | Steam | Magnesium oxide and hydrogen | |
All group 2 elements | Oxygen | Metal oxide | |
Strontium and barium | Oxygen | Metal peroxide | |
All group 2 elements | Chlorine | Metal chloride | |
All group 2 elements | Hydrochloric acid | Metal chloride and hydrogen | |
All group 2 elements | Nitric acid | Metal nitrate and hydrogen | |
All group 2 elements | Sulphuric acid | Metal sulphate and hydrogen |
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 putting them in a flame. The different metals burn to produce different-coloured flames in a spectacular show.
Get a clean metal loop and dip it in acid. Hold it 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 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.
Lastly, let's focus on some of the uses of group 2.
Check out Group 2 Compounds for more uses of group 2.
Group 2 on the periodic table refers to the alkaline earth metals.
Group 2 elements are fairly reactive metals with relatively low densities and melting and boiling points. Their reactivity and atomic radius increase as you move down the group, whilst their melting point, electronegativity, and first ionisation energy decrease as you move down the group.
Group 2 elements are called alkaline earth metals because they form alkaline metal oxides and hydroxides. These metal oxides are found in the Earth.
The solubility of group 2 hydroxides increases as you move down the group, whereas the solubility of group 2 sulphates decreases as you move down the group.
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