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?
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Jetzt kostenlos anmeldenGroup 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?
Firstly, let's answer one basic question: what are group 2 compounds?
Group 2 compounds are ionic compounds containing a group 2 metal cation.
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.
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:
Check out the article Group 2 for more information about group 2 metals. In the explanation, you'll also be able to learn about the properties and uses of these elements.
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:
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.
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:
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:
Later on in the article, you'll find out more about how the differing solubilities of group 2 compounds contribute to their everyday uses.
Solubility depends on the compound's enthalpy change of solution (ΔHs°). 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 (ΔHL°), and the total enthalpy change of hydration (ΔHh°) 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:
Similar principles apply to the solubility of group 2 hydroxides, but with a different outcome:
Enthalpy of Solution and Hydration will explore factors affecting the enthalpy change of solution in more detail.
Next up, let's find out about the thermal stabilities of group 2 compounds.
Luckily for you, you have to remember just one trend for the thermal stability of both group 2 nitrates and carbonates:
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):
$$2Mg(NO_3)_2(s)\rightarrow 2MgO(s)+4NO_2(g)+O_2(g)$$
$$MgCO_3(s)\rightarrow MgO(s)+CO_2(g)$$
NO2 is toxic, so the thermal decomposition of group 2 nitrates must always be carried out in a fume cupboard.
Once again, here's a table summarising the new information:
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.
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:
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.
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:
Here's a general equation for the reaction of a group 2 oxide with water:
$$MO(s)+H_2O(l)\rightarrow M(OH)_2(aq)$$
Group 2 oxides, hydroxides, and carbonates, all react with dilute hydrochloric acid to form a chloride salt (MCl2), along with other products:
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:
$$M(OH)_2(aq)+2HCl(aq)\rightarrow MCl_2(aq)+2H_2O(l)$$
Remember that not all group 2 hydroxides are soluble - some could be solid instead!
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.
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.
Again, remember that group 2 hydroxides vary in solubility - some might be solid in solution, whilst others are aqueous. However, group 2 oxides and carbonates are all insoluble.
Here's how group 2 carbonates react with sulfuric acid to produce a soluble sulfate:
$$MCO3(S)+H_2SO_4(aq)\rightarrow MSO_4(aq)+H_2O(l)+CO_2(g)$$
To help consolidate your learning, we've made a handy table bringing together the various reactions of group 2 compounds:
Group 2 compound | Water | Hydrochloric acid | Sulfuric acid |
Oxide | Hydroxide (dissolves with varying solubility) | Chloride salt + water | Sulfate salt + water |
Hydroxide | Dissolves with varying solubility | Chloride salt + water | Sulfate salt + water |
Carbonate | Insoluble (no reaction) | Chloride salt + water + carbon dioxide | Sulfate salt + water + carbon dioxide |
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:
Check out Test-Tube Reactions for more ways of identifying unknown ions.
A group 2 compound is a compound containing a group 2 metal cation with a charge of 2+.
Properties of group 2 compounds include their solubility and thermal stability:
To identify group 2 compounds, you need to test for both the positive group 2 metal cation, and the remaining negative anion(s):
Check out our article Test-Tube Reactions for more information on the tests for inorganic ions.
Examples of group 2 compounds include calcium carbonate (CaCO3), barium sulfate (BaSO4), and magnesium hydroxide (Mg(OH)2).
Some group 2 compounds are soluble and form colourless solutions. However, some group 2 compounds are insoluble. These are generally found as white solids.
What are group 2 compounds?
Compounds that contain a group 2 metal cation with a charge of +2.
The solubility of group 2 sulfates ____ as you move down the group in the periodic table.
Decreases.
Which is more soluble?
CaSO4
Group 2 nitrates become ____ thermally stable as you move down the group in the periodic table.
More
What is Ca(OH)2 used for?
In agriculture, to increase the pH of soils. It is also known as limewater.
What is Mg(OH)2 used for?
To treat indigestion. It is also known as milk of magnesia.
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