Group 2 Reactivity

Did you know that magnesium, a Group 2 metal, is the fourth most common element found in the Earth itself, and the third most common ion in seawater? It is essential to human life too; magnesium plays a role in the structures of 300 different enzymes. But we never find magnesium, or in fact any of the other group 2 metals, in their elemental states in nature. This is because they always react with other species to form ions and compounds. In this article, we'll explore these reactions and look more closely at Group 2 reactivity.

• This article is about group 2 reactivity in inorganic chemistry.
• We'll define group 2 and briefly look at their electron configurations.
• We'll then dive into the Redox reaction of group 2.
• This will involve their reactions with water (to produce hydroxides), oxygen, chlorine, and acids.
• We'll also find out about a further reaction of magnesium, used to extract titanium industrially.
• We'll finish by summarising all that we've learned in a handy table before

Group 2 definition

Group 2 contains the following metals:

• Beryllium (Be).
• Magnesium (Mg).
• Calcium (Ca).
• Strontium (Sr).
• Barium (Ba).
• Radium (Ra).

Group 2 metals are found in the s-block in the Periodic Table. They all have two electrons in their outer shell and, except for beryllium, form positive cations with a charge of 2+. For example, magnesium has the electron configuration 1s² 2s² 2p⁶ 3s², and loses two electrons to form Mg²⁺ ions with the configuration 1s² 2s² 2p⁶.

Fig. 1: The position of group 2 (highlighted in pink) in the periodic table.StudySmarter Originals

Group 2 reactivity: redox reactions

Now that we've learned the basics about group 2, we can take a look at their reactions and reactivity. In fact, all reactions involving group 2 metals are Redox Reactions.

The group 2 element loses two electrons to form a cation with charge 2+, and an oxidation state of +2. It is oxidised. On the other hand, the other species involved gains electrons. It is reduced. Don't worry - we'll practice showing the processes of oxidation and reduction below.

Let's now explore the reactions of group 2 with water, oxygen, chlorine, and acids.

Group 2 reactivity with water

Have you ever wondered why group 2 elements are called the alkaline metals? It's because they react with water to produce alkaline metal hydroxides. This is a great example of the trend in reactivity of group 2 elements.

Reaction

As we said, group 2 metals (which we've represented using M) react with water (H₂O) to produce a metal hydroxide (M(OH)₂). This reaction also releases hydrogen gas (H₂). Here's the general equation:

$$M(s)+2H_2O(l)\rightarrow M(OH)_2(aq/s)+H_2(g)$$

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.

Fig. 2: The oxidation states in magnesium hydroxide. StudySmarter Originals

Reactivity

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. In fact, the overall reactivity of the alkaline metals increases as you move down the group. This is because their ionisation energies decrease as you go down the group. A lower Ionisation Energy means that it is easier to lose electrons, and so reactivity increases.

Group 2 reactivity with oxygen

Our next reaction of interest is the one between group 2 elements and oxygen.

Reaction

Burning a group 2 element in oxygen (O₂) produces a metal oxide (MO). 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 (MO₂). However, beryllium, magnesium, and calcium don't. This has to do with the smaller size of these metal ions. Since Be²⁺, Mg²⁺, and Ca²⁺are smaller ions than St²⁺ and Ba²⁺, they have a much higher charge density, and so attract and bond with oxygen molecules differently. Here's the equation for the formation of strontium or barium perioxide:

$$M(s)+O_2(g)\rightarrow MO_2(s)$$

Reactivity

It is hard to give clear trends when it comes to the reactivity of group 2 metals with oxygen. Their reactivity depends on many factors, such as the heat needed to start the reaction and whether the metal is covered with a protective layer of metal oxide or not.

Group 2 reactivity with chlorine

Next up: the reaction between group 2 elements and chlorine.

Reaction

Group 2 metals react with chlorine (Cl₂) to produce metal chlorides (MCl₂). In fact, group 2 metals react similarly with all halides. Here's the general equation:

$$M(s)+Cl_2(g)\rightarrow MCl_2(s)$$

Reactivity

The reactivity between group 2 elements and chlorine increases as you go down the group, thanks to the decreasing ionisation energy of the group 2 metal.

Group 2 reactivity with acids

Finally, we'll turn our attention to the reactions of group 2 with acids.

Reaction

The reactions between group 2 metals and acids vary depending on the type of acid used, but all form colourless solutions of ionic salts.

Hydrochloric acid

This is the simplest case - group 2 elements react with hydrochloric acid (HCl) to give a metal chloride (MCl₂) and hydrogen gas (H₂):

$$M+2HCl(aq)\rightarrow MCl_2(aq)+H_2(g)$$

As before, the metal is oxidised, whilst hydrogen is reduced. The chloride ions are spectator ions; they are neither oxidised nor reduced, but rather keep the same physical and oxidation state.

Nitric acid

The reaction with nitric acid is a little more complicated. Remember that group 2 metal reactions are Redox Reactions. Here, the group 2 metal is oxidised, and hydrogen ions are reduced. But the group 2 metal can also reduce the Nitrogen atoms found within the nitrate ion (NO₃^-), forming all manner of nitrous oxides. At A level, we ignore these Nitrous Oxide products and instead just consider the main products: a metal nitrate (M(NO₃)₂) and hydrogen gas (H₂). Like usual, we've given you the general equation:

$$M(s)+2HNO_3(aq)\rightarrow M(NO_3)_2(aq)+H_2(g)$$

Sulfuric acid

The last group 2 metal-acid reaction we'll look it is their reaction with sulfuric acid (H₂SO₄). In theory, group 2 elements react with sulfuric acid to produce metal sulfates (MSO₄) and hydrogen gas (H₂). However, in reality, strontium and barium don't tend to react much in this way. This is because strontium and barium sulfate (SrSO₄ and BaSO₄) are both insoluble, and so the sulfate forms an impenetrable layer on the surface of the metal and prevents any further reaction. Likewise, calcium sulfate (CaSO₄) is only sparingly soluble, so calcium doesn't react much with sulfuric acid either. Nevertheless, here's the general reaction between a group 2 metal and sulfuric acid:

$$M(s)+H_2SO_4(aq)\rightarrowMSO_4(aq)+H_2(g)$$

Now practice applying these three general equations to a specific group 2 metal.

Further reactions of group 2

One further reaction involving group 2 metals uses magnesium to extract titanium (Ti). 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 (TiCl₄). Magnesium is then added. This reduces the titanium within titanium tetrachloride, forming magnesium chloride (MgCl₂) and titanium metal (Ti):

$$TiCl_4(l)+2Mg(s)\rightarrow 2MgCl_2(s)+Ti(s)$$

Summary of group 2 reactivity

Here's a handy table that summarises the reactions of group 2 for you. For each reaction, we've included the relevant element(s) in group 2, the other reactant(s), the product(s), and the general equation.