Group 2 Compounds

Group 2 metals, also known as alkaline earth metals, have the ability to form a variety of compounds with a wide range of uses. From chemical analysis to medical diagnosis, these compounds have several roles in daily life. But how do these compounds form, how do they react, and why do their chemical properties make them suitable for such uses?

• This article is about group 2 compounds in inorganic chemistry.
• Firstly, we will look at the meaning of group 2 compounds and name some examples.
• We'll then dive into their properties, including trends in solubility and thermal stability.
• Subsequently, we shall explore the reactions of group 2 compounds with water and dilute acids.
• Finally, we will outline some of the uses of group 2 compounds.

Group 2 compounds: meaning and examples

Firstly, let's answer one basic question: what are group 2 compounds?

Group 2 elements are also known as the alkaline earth metals and are a part of the s-block on the periodic table. They all have two electrons in their outermost shell.

Fig. 1: A diagram of the periodic table, highlighting the group 2 metals. StudySmarter Originals

When group 2 metal atoms react to form ions, they lose their two outer electrons and so form positive cations with a charge of +2. These cations can bond with a range of negative anions, forming compounds with a myriad of different properties and uses. The negative anions in group 2 compounds always have a combined total charge of -2.

Here are examples of group 2 compounds. We've used the letter M to represent a general group 2 metal:

• Hydroxides, M(OH)₂.
• Carbonates, MCO₃.
• Sulfates, MSO₄.
• Nitrates, M(NO₃)₂.

Properties of group 2 compounds

So, we now know what group 2 compounds are. But how do they behave, and what are their differences? We'll explore this by considering their properties and reactions. To start, we'll look at the properties of group 2 compounds, including:

• Solubility.
• Thermal stability.

Solubility

If you've ever had an x-ray of your digestive tract, you probably ate a meal of barium sulfate (BaSO₄) beforehand. This white group 2 compound is insoluble in water and so shows up in the x-ray, helping outline the features of your gut. You couldn't use magnesium sulfate (MgSO₄), another group 2 compound, for this purpose - it is much more soluble in water. Instead, we use magnesium sulfate in bath salts and marine aquaria, to increase concentrations of aqueous Mg²⁺. We'll now investigate the solubility of group 2 compounds, including the observable trends and their explanation.

Trend

You might be able to guess from the information above that group 2 sulfates become less soluble as you move down the group in the periodic table. However, other group 2 compounds show different trends in solubility:

• As we mentioned, the solubility of group 2 sulfates decreases as you go down the group in the periodic table.
• On the other hand, the solubility of group 2 hydroxides increases as you move down the group.

For example, here's an equation representing the dissolution of a soluble group 2 hydroxide, barium hydroxide:

$$Ba(OH)_2(s)\rightarrow Ba^{2+}(aq)+2OH^-(aq)$$

We've summarised this information into a handy table for you:

Fig. 2: A table showing the trend in solubility of group 2 hydroxides and sulfates.StudySmarter Originals

Explanation

Solubility depends on the compound's enthalpy change of solution (ΔH_s°). The more positive (more endothermic) the enthalpy change, the less soluble the compound. Enthalpy change of solution is in turn affected by the compound's lattice enthalpy (ΔH_L°), and the total enthalpy change of hydration (ΔH_h°) of the ions within the compound. The greater the magnitude of the lattice enthalpy, and the smaller the magnitude of the combined enthalpy changes of hydration, the more positive the overall enthalpy change of solution, and so the less soluble the compound.

Let's consider the enthalpy of solution of group 2 sulfates:

• Lattice enthalpy decreases in magnitude as you move down group 2 in the periodic table. This is because the size of the group 2 metal cation increases, but its charge stays the same, reducing its overall charge density and its attraction to the sulfate anion.
• The group 2 metal cation's enthalpy of hydration also decreases in magnitude as you move down the group. This is again due to the relative size of the ions.
• However, lattice enthalpy decreases in relatively small steps. This is because of the large size of the sulfate ions compoared to the size of themetal cations. On the other hand, enthalpy of hydration decreases in relatively large steps. This means that overall, the enthalpy change of solution increases(becomes more endothermic) as you move down the group in the periodic table. Thus, solubility increases.

Fig. 3: An enthalpy diagram explaining the trend in solubilty of group 2 sulfates.StudySmarter Originals

Similar principles apply to the solubility of group 2 hydroxides, but with a different outcome:

• Lattice enthalpy again decreases in magnitude as you move down the group in the periodic table.
• The combined enthalpy of hydration also decreases in magnitude as you move down the group.
• However, lattice enthalpies decrease in larger steps than the combined enthalpies of hydration due to the relatively small size of the hydroxide ion compared to the positive metal cation. This means that overall, the enthalpy change of solution decreases (becomes less endothermic) as you move down the group in the periodic table. Thus, solubility increases.

Fig. 4: An enthalpy diagram explaining the trend in solubility of group 2 hydroxides.StudySmarter Originals

Thermal stability

Next up, let's find out about the thermal stabilities of group 2 compounds.

Trend

Luckily for you, you have to remember just one trend for the thermal stability of both group 2 nitrates and carbonates:

• The thermal stability of group 2 nitrates increases as you move down the group in the periodic table.
• Likewise, the thermal stability of group 2 carbonates also increases as you move down the group.

This means that group 2 nitrates and carbonates require heating to higher temperatures before they decompose, as you move down from magnesium to barium in the periodic table. Group 2 nitrates thermally decompose into metal oxides (MO), nitrogen dioxide (NO₂), and oxygen (O₂), whilst group 2 carbonates thermally decompose into just metal oxides (MO) and carbon dioxide (CO₂).

For example, here are equations showing the thermal decomposition of magnesium nitrate (Mg(NO₃)₂) and magensium carbonate (MgCO₃):

$$2Mg(NO_3)_2(s)\rightarrow 2MgO(s)+4NO_2(g)+O_2(g)$$

$$MgCO_3(s)\rightarrow MgO(s)+CO_2(g)$$

Once again, here's a table summarising the new information:

Fig. 5: A table summarising the trends in thermal stability of group 2 nitrates and carbonates.StudySmarter Originals

Explanation

Thermal stability depends on the enthalpy change of the decomposition reaction (ΔH°_r). The more positive the enthalpy change, the more thermally stable the compound, and the higher the temperature it needs to thermally decompose. But what causes differences in the enthalpy change of the decomposition reaction? It is all down to the size of the group 2 cation and its relative polarising ability. We'll look at this in terms of group 2 carbonates, but the same ideas apply to group 2 nitrates.

• When carbonates thermally decompose, one of the C-O bonds in the CO₃^2- carbonate ion breaks, releasing CO₂.
• As you move down group 2 in the periodic table, the size of the metal cations increases.
• However, all group 2 metal cations have the same charge. This means that as you move down the group, charge density decreases. As a result, the ion's polarising ability decreases.
• Smaller metal ions are much better at polarising other species than larger metal cations. Thus, smaller group 2 cations are more able to polarise the carbonate ion than larger group 2 cations, weakening one of the carbonate ion's C-O bonds. This makes the bond easier to break.
• As a result, the carbonate ion thermally decomposes much more easily.
• Therefore, group 2 metal carbonates with smaller metal cations are much less thermally stable than those with larger metal cations, and so thermally stability increases as you move down the group.

Reactions of group 2 compounds

Do you know of the applications of the group 2 compound calcium chloride (CaCl₂)? For example, even in just the food industry, it is used in a variety of ways:

• To help firm tinned fruit and tofu, a common plant-based protein source.
• In sports drinks as an electrolyte.
• Thanks to its salty taste, it is also used to flavour pickles without increasing their sodium content.

Calcium chloride is made by reacting group 2 compounds with dilute hydrochloric acid. Let's now explore this reaction further, alongside other reactions of group 2 compounds. We'll consider how group 2 oxides (MO), hydroxides (M(OH)₂), and carbonates (MCO₃), react with water, hydrochloric acid, and sulfuric acid.

Reaction with water

We've already considered how group 2 hydroxides react with water - they dissolve with varying solubility. But we haven't yet seen the reactions between group 2 oxides or carbonates with water:

• Group 2 oxides react with water to form a hydroxide, M(OH)₂.
• Group 2 hydroxides dissolve in water with varying solubility. Solubility increases as you move down the group in the periodic table. The solutions also become increasingly alkaline, which gives the alternate name for group 2 elements: the alkaline metals.
• On the other hand, group 2 carbonates are insoluble in water - they don't react at all.

Here's a general equation for the reaction of a group 2 oxide with water:

$$MO(s)+H_2O(l)\rightarrow M(OH)_2(aq)$$

Reaction with hydrochloric acid

Group 2 oxides, hydroxides, and carbonates, all react with dilute hydrochloric acid to form a chloride salt (MCl₂), along with other products:

• Group 2 oxides react with dilute hydrochloric acid to form a chloride salt and water.
• Group 2 hydroxides also react with dilute hydrochloric acid to form a chloride salt and water.
• Group 2 carbonates react with dilute hydrochloric acid to form a chloride salt, water, and carbon dioxide.

Group 2 chlorides are highly soluble, and so all the compounds dissolve readily in solution. Here's the general equation for the reaction of a group 2 hydroxide with hydrochloric acid:

$$M(OH)_2(aq)+2HCl(aq)\rightarrow MCl_2(aq)+2H_2O(l)$$

Reaction with sulfuric acid

Last up: how do group 2 oxides, hydroxides, and carbonates react with dilute sulfuric acid? It's simple - they produce sulfate salts. Once again, the additional products vary.

• Group 2 oxides react with dilute sulfuric acid to form a sulfate salt and water.
• Group 2 hydroxides also react with dilute sulfuric acid to form a sulfate salt and water.
• Group 2 carbonates react with dilute sulfuric acid to form a sulfate salt, water, and carbon dioxide.

The extent of the reaction depends on the solubility of the sulfate formed and whether the reactant is solid or not. If the metal sulfate is insoluble, it precipitates out of solution and onto the surface of any solid reactant. This prevents any further reaction from taking place. On the other hand, if the sulfate formed is soluble, the reaction continues on.

Here's how group 2 carbonates react with sulfuric acid to produce a soluble sulfate:

$$MCO₃(S)+H_2SO_4(aq)\rightarrow MSO_4(aq)+H_2O(l)+CO_2(g)$$

Summary

To help consolidate your learning, we've made a handy table bringing together the various reactions of group 2 compounds:

Uses of Group 2 compounds

To round off the article, let’s talk about some uses of common group 2 compounds. However, remember that this list is not exhaustive - group 2 compounds have hundreds of different applications! They span all sorts of industries, from healthcare and pharmaceuticals to agriculture and construction:

• Magnesium oxide (MgO) is a white solid and can be used as a heat-resistant ceramic to line furnaces due to its very high melting temperature.
• MgO turns to magnesium hydroxide (Mg(OH)₂) when reacted with water. This insoluble compound is the active ingredient in milk of magnesia, used to help with indigestion.
• Calcium hydroxide (Ca(OH)₂) is normally known as limewater. It is used in agriculture to increase the pH of soil.
• On the other hand, calcium oxide (CaO) can be used to remove sulfur dioxide (SO₂) pollutants from flue gases. SO₂ is formed when fossil fuels are burnt to produce electricity.
• Barium sulfate (BaSO₄) is used in medicine because of its insolubility. It absorbs X-rays strongly and is used to diagnose disorders of the intestines and stomach. Because it is insoluble, it is not absorbed into the bloodstream from the gut.
• Acidified barium chloride (BaCl₂) is used to test for sulfate ions. This is done by adding dilute hydrochloric acid to the unknown solution, followed by barium chloride. If a white precipitate of insoluble barium sulfate forms, we know that the sulfate ions are present.