The Chemistry of Group 2: The Alkaline Earth Metals
Welcome to the exciting world of Inorganic Chemistry! This chapter focuses on Group 2 of the Periodic Table, commonly known as the Alkaline Earth Metals. These elements—Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), and Barium (Ba)—are highly reactive and play crucial roles, from medicine to industry.
Understanding Group 2 is essential because it demonstrates fundamental trends in atomic structure and reactivity that apply across the whole Periodic Table. Don't worry if periodic trends seem tricky; we’ll break them down using simple explanations!
3.2.2.1 Trends in Physical Properties (Mg to Ba)
As you move down Group 2, the elements show clear and predictable trends in their atomic properties. These changes are all rooted in the increasing number of electron shells.
Trend 1: Atomic Radius
Observation: Atomic radius increases down the group (Mg to Ba).
Explanation:
- Moving down the group, each element has one more complete electron shell than the element above it.
- Although the nuclear charge (number of protons) increases, the outer electrons are further from the nucleus.
- The increasing number of inner shells provides greater shielding, reducing the effective pull of the nucleus on the outer electrons.
- Analogy: Think of wrapping layers of clothes (shells) around a core. The outermost layer gets pushed further and further out.
Trend 2: First Ionisation Energy (IE)
Definition Recap: The first ionisation energy is the minimum energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.
Observation: First ionisation energy decreases down the group (Mg to Ba).
Explanation:
- The atoms get larger (Trend 1), meaning the outermost electron is further away from the nucleus.
- The shielding effect from the inner electrons increases.
- Therefore, the electrostatic attraction between the nucleus and the outer electron is weaker, making it easier to remove the electron.
Key Takeaway: Because it gets easier to lose electrons down the group, reactivity generally increases from Mg to Ba.
Trend 3: Melting Points
Structure and Bonding: Group 2 metals exist as giant metallic lattices. They consist of positive ions (\( \text{M}^{2+} \)) surrounded by a sea of delocalised electrons.
Observation: Melting points generally decrease down the group (Mg to Ba), although Magnesium is often an exception (it has a relatively high MP).
Explanation (Structure and Bonding):
- The metallic bond strength determines the melting point.
- All Group 2 elements have a constant ionic charge (2+) and the same number of delocalised electrons per atom (two).
- As you move down the group, the size of the ion (\( \text{M}^{2+} \)) increases.
- The delocalised electrons are spread over a larger volume, meaning the electrostatic attraction between the positive ions and the sea of electrons is weaker.
- Weaker metallic bonds require less energy to break, leading to a lower melting point.
Quick Review: Trends Down Group 2
- Atomic Radius: Increases (More shells/more shielding).
- First Ionisation Energy: Decreases (Easier to lose electrons).
- Reactivity: Increases.
- Melting Point: Generally Decreases (Weaker metallic bonding).
Reactions of Group 2 Elements
1. Reaction with Water
Group 2 metals react with water to form the metal hydroxide and hydrogen gas. The vigour of this reaction increases down the group, matching the trend in reactivity.
Magnesium (Mg):
Magnesium reacts very slowly with cold water, as a layer of insoluble \( \text{Mg(OH)}_2 \) quickly forms on the surface, stopping the reaction (passivation). However, it reacts much more vigorously with steam:
\( \text{Mg(s)} + \text{H}_2\text{O(g)} \rightarrow \text{MgO(s)} + \text{H}_2\text{(g)} \)
(Magnesium oxide is formed instead of the hydroxide because \( \text{Mg(OH)}_2 \) decomposes at high temperatures.)
Calcium (Ca), Strontium (Sr), Barium (Ba):
These react increasingly vigorously with cold water to form the metal hydroxide and hydrogen gas. Calcium, for example, produces a white suspension (as \( \text{Ca(OH)}_2 \) is sparingly soluble, see below):
\( \text{Ca(s)} + 2\text{H}_2\text{O(l)} \rightarrow \text{Ca(OH)}_2\text{(aq/s)} + \text{H}_2\text{(g)} \)
2. Industrial Use: Magnesium in Titanium Extraction
Magnesium is vital in extracting transition metals like titanium (Ti). Titanium cannot be extracted economically using carbon (like iron) because it forms titanium carbide, so a more reactive metal is needed.
Magnesium acts as a powerful reducing agent in the high-temperature reaction to convert titanium(IV) chloride into titanium metal:
\( \text{TiCl}_4\text{(g)} + 2\text{Mg(l)} \rightarrow \text{Ti(s)} + 2\text{MgCl}_2\text{(l)} \)
Did you know? Titanium is used in aircraft because it is strong, light, and corrosion-resistant.
Key Takeaway: Reactivity with water increases down Group 2, and magnesium's reducing power makes it useful for extracting titanium.
Solubility Trends and Applications (Crucial AS Content!)
The solubility of Group 2 compounds is perhaps the most important (and sometimes confusing) part of this section. You need to know two distinct trends: the solubility of hydroxides and the solubility of sulfates.
1. Solubility of Group 2 Hydroxides (M(OH)₂)
Observation: The solubility of Group 2 hydroxides increases down the group (Mg to Ba).
Memory Aid: "OH! Solubility is going UP!"
- \( \text{Mg(OH)}_2 \) (Magnesium Hydroxide) is sparingly soluble (or virtually insoluble).
- \( \text{Ba(OH)}_2 \) (Barium Hydroxide) is very soluble.
Uses of Hydroxides:
- Magnesium Hydroxide, \( \text{Mg(OH)}_2 \): Used in medicine as an antacid (like Milk of Magnesia) to treat indigestion. It neutralises excess stomach acid (\( \text{HCl} \)):
\( \text{Mg(OH)}_2\text{(s)} + 2\text{HCl(aq)} \rightarrow \text{MgCl}_2\text{(aq)} + 2\text{H}_2\text{O(l)} \)Because it is only sparingly soluble, it is not harsh or toxic.
- Calcium Hydroxide, \( \text{Ca(OH)}_2 \): Known as 'slaked lime', it is used in agriculture to change soil pH (neutralise acidic soil). Acidic soil reduces crop yields, so adding lime helps maintain the food supply.
2. Solubility of Group 2 Sulfates (MSO₄)
Observation: The solubility of Group 2 sulfates decreases down the group (Mg to Ba).
Memory Aid: "Sulfates Sink! Down the group, they are Insoluble."
- \( \text{MgSO}_4 \) (Magnesium Sulfate) is soluble.
- \( \text{BaSO}_4 \) (Barium Sulfate) is insoluble.
Uses of Sulfates:
- Barium Sulfate, \( \text{BaSO}_4 \): Used in medicine as a barium meal. Patients swallow it before an X-ray of the digestive system. Because \( \text{BaSO}_4 \) is highly insoluble, it is safe despite Barium ions being toxic. It blocks X-rays, allowing doctors to see the stomach and intestines clearly.
Accessibility Tip: Solubilities Swap!
For Group 2 elements, when you go down the group:
Hydroxides (\( \text{OH} \)) = More Soluble (Up)
Sulfates (\( \text{SO}_4 \)) = Less Soluble (Down)
3. Industrial and Analytical Chemistry
1. Flue Gas Desulfurisation (FGD)
When fossil fuels containing sulfur impurities are burned, they produce sulfur dioxide (\( \text{SO}_2 \)), which causes acid rain. Group 2 compounds are used to remove this pollutant from industrial chimney gases (flue gases).
We use basic Group 2 compounds, typically Calcium Oxide (\( \text{CaO} \)) or Calcium Carbonate (\( \text{CaCO}_3 \)), which act as bases to neutralise the acidic sulfur dioxide.
Reaction using Calcium Carbonate:
\( \text{CaCO}_3\text{(s)} + \text{SO}_2\text{(g)} \rightarrow \text{CaSO}_3\text{(s)} + \text{CO}_2\text{(g)} \)
Reaction using Calcium Oxide:
\( \text{CaO(s)} + \text{SO}_2\text{(g)} \rightarrow \text{CaSO}_3\text{(s)} \)
This process reduces environmental harm by preventing acid rain formation.
2. Testing for Sulfate Ions (\( \text{SO}_4^{2-} \))
Because Barium Sulfate (\( \text{BaSO}_4 \)) is famously insoluble, we use a solution containing Barium ions to test for sulfates.
Reagent: Acidified Barium Chloride solution (\( \text{BaCl}_2\text{(aq)} \)).
Process:
- Add a few drops of dilute hydrochloric acid (\( \text{HCl} \)) or nitric acid to the sample solution.
- Add the \( \text{BaCl}_2\text{(aq)} \) solution.
- If sulfate ions are present, a white precipitate of barium sulfate is immediately formed.
Ionic Equation:
\( \text{Ba}^{2+}\text{(aq)} + \text{SO}_4^{2-}\text{(aq)} \rightarrow \text{BaSO}_4\text{(s)} \)
Why must the \( \text{BaCl}_2 \) solution be acidified?
This is a common exam question! The solution must be acidified (usually with \( \text{HCl} \)) to react with and remove other ions that also form insoluble white precipitates with barium ions, such as carbonate ions (\( \text{CO}_3^{2-} \)) or sulfite ions (\( \text{SO}_3^{2-} \)).
If these interfering ions were present, they would give a false positive result (a white ppt that isn't \( \text{BaSO}_4 \)). Acidifying the solution converts these ions into soluble compounds or gases (e.g., \( \text{CO}_2 \)), preventing them from confusing the test.
Key Takeaway: Group 2 chemistry is fundamental to environmental protection (FGD) and analytical chemistry (the sulfate test).