Group II elements and their compounds

Q: What is the general trend in first and second ionisation energies down Group 2, and why?

Ionisation energies decrease down the group because atomic radius and electron shielding increase, which outweighs the increase in nuclear charge and weakens the attraction on outer electrons.

Q: Why do melting points generally decrease down Group 2?

Down the group, the ionic radius of the \(\text{M}^{2+}\) cation increases while the charge stays constant, decreasing the charge density and weakening the metallic bonding with delocalised electrons.

Q: What are the different products when magnesium reacts with cold water versus steam?

With cold water, magnesium reacts slowly to form magnesium hydroxide (\(\text{Mg(OH)}_2\)) and \(\text{H}_2\); with steam, it reacts vigorously to form magnesium oxide (\(\text{MgO}\)) and \(\text{H}_2\).

Q: What are the solubility trends for Group 2 hydroxides and sulfates down the group?

Hydroxide solubility increases down the group, whereas sulfate solubility decreases down the group.

Q: Why does the thermal stability of Group 2 carbonates increase down the group?

Larger cations down the group have lower charge density and lower polarising power, causing less distortion of the carbonate ion's \(\text{C–O}\) bonds.

Q: What are the products and observations during the thermal decomposition of Group 2 nitrates?

They decompose into a metal oxide, brown nitrogen dioxide gas (\(\text{NO}_2\)), and oxygen gas (\(\text{O}_2\)).

Q: What flame colours are produced by \(\text{Ca}^{2+}\), \(\text{Sr}^{2+}\), \(\text{Ba}^{2+}\), and \(\text{Mg}^{2+}\)?

\(\text{Ca}^{2+}\) gives brick-red, \(\text{Sr}^{2+}\) gives crimson-red, \(\text{Ba}^{2+}\) gives apple-green, and \(\text{Mg}^{2+}\) gives no flame colour.

Q: How do you test for sulfate ions (\(\text{SO}_4^{2-}\)) in solution?

Acidify the sample with dilute \(\text{HCl}\) (or \(\text{HNO}_3\)), then add aqueous \(\text{BaCl}_2\); a white precipitate of \(\text{BaSO}_4\) confirms sulfate ions.

Q: Why is barium sulfate (\(\text{BaSO}_4\)) safe to ingest for medical X-ray imaging despite \(\text{Ba}^{2+}\) ions being toxic?

\(\text{BaSO}_4\) is completely insoluble in water, preventing toxic \(\text{Ba}^{2+}\) ions from being absorbed into the bloodstream.

Chemistry 1110 · CCEA AS Level · 9 min read · Updated: Aug 31, 2026

Introduction to Group 2: The Alkaline Earth Metals

Welcome to your study notes for Group 2 of the Periodic Table! Often called the alkaline earth metals, this group includes beryllium (\(\text{Be}\)), magnesium (\(\text{Mg}\)), calcium (\(\text{Ca}\)), strontium (\(\text{Sr}\)), and barium (\(\text{Ba}\)).

These elements are all around us in daily life: calcium gives strength to our bones and teeth, magnesium is at the heart of chlorophyll in green plants, and barium compounds help doctors take detailed X-ray images of the digestive tract. In this chapter, we will explore their physical trends, how they react with oxygen and water, the solubility and thermal stability of their compounds, and how to identify them using simple chemical tests.

Don't worry if periodic trends felt overwhelming in previous units. Group 2 follows logical patterns that make sense once you understand the underlying atomic structure!

Key Takeaway: All Group 2 elements have two outer-shell electrons (\(s^2\)) and form \(2+\) ions (\(\text{M}^{2+}\)) by losing these two electrons during chemical reactions.

1. Atomic and Physical Trends Down Group 2

As you move down Group 2 from \(\text{Mg}\) to \(\text{Ba}\), several fundamental atomic properties change in predictable ways.

a) Electron Configuration

Every Group 2 element has two electrons in its highest energy \(s\)-subshell:

• Magnesium (\(\text{Mg}\)): \([\text{Ne}]\,3s^2\)
• Calcium (\(\text{Ca}\)): \([\text{Ar}]\,4s^2\)
• Strontium (\(\text{Sr}\)): \([\text{Kr}]\,5s^2\)
• Barium (\(\text{Ba}\)): \([\text{Xe}]\,6s^2\)

b) Atomic Radius

Trend: Atomic radius increases down the group.
Explanation: Each successive element down the group has an extra quantum shell of electrons. This additional electron shell increases the physical distance between the outer electrons and the nucleus, increasing the size of the atom despite the increasing nuclear charge.

c) First and Second Ionisation Energies

Trend: Ionisation energies decrease down the group.
Explanation: Recall that first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms: \(\text{M(g)} \rightarrow \text{M}^+\text{(g)} + \text{e}^-\). As we go down the group:
1. The atomic radius increases, meaning outer electrons are further from the positive nucleus.
2. There is more electron shielding from inner complete electron shells.
3. Even though nuclear charge increases (more protons), the increased shielding and distance outweigh this, leading to a weaker electrostatic attraction on the outermost electrons.
4. Therefore, less energy is needed to remove the first and second electrons.

d) Melting Points

Trend: Melting points generally decrease down the group (with \(\text{Mg}\) being an anomaly having a slightly lower melting point due to a difference in crystal packing structure).
Explanation: Group 2 elements form metallic lattices containing \(2+\) metal ions surrounded by a "sea" of delocalised electrons. Down the group, the ionic radius of the metal cation increases, while the number of delocalised electrons per ion remains constant (+2). This means the charge density decreases, leading to a weaker electrostatic attraction between the larger cations and the delocalised electrons.

Quick Review: Down Group 2 \(\rightarrow\) Atomic radius increases \(\rightarrow\) Ionisation energy decreases \(\rightarrow\) Reactivity increases.

2. Chemical Reactivity of Group 2 Elements

Group 2 metals act as powerful reducing agents because they readily donate their two outer electrons to form \(\text{M}^{2+}\) ions, undergoing oxidation: \(\text{M} \rightarrow \text{M}^{2+} + 2\text{e}^-\).

Reactivity Trend: Reactivity with both oxygen and water increases down the group because it becomes progressively easier to remove the two outer electrons.

a) Reaction with Oxygen

All Group 2 metals react with oxygen when heated in air to form standard solid white ionic metal oxides with the general formula \(\text{MO}\):

\(2\text{M(s)} + \text{O}_2\text{(g)} \rightarrow 2\text{MO(s)}\)

Example: When magnesium burns in air, it produces a dazzling, intense white flame and leaves behind a white solid powder of magnesium oxide:

\(2\text{Mg(s)} + \text{O}_2\text{(g)} \rightarrow 2\text{MgO(s)}\)

b) Reaction with Water

The reaction of Group 2 metals with water highlights the clear increase in reactivity down the group:

• Magnesium with Cold Water: Reacts very slowly over days. It forms a weakly alkaline suspension of magnesium hydroxide and bubbles of hydrogen gas:
\(\text{Mg(s)} + 2\text{H}_2\text{O(l)} \rightarrow \text{Mg(OH)}_2\text{(s)} + \text{H}_2\text{(g)}\)

• Magnesium with Steam (Crucial Exam Reaction!): Reacts vigorously and rapidly. Heated magnesium burns with a bright white flame in steam to form white solid magnesium oxide (not hydroxide) and hydrogen gas:
\(\text{Mg(s)} + \text{H}_2\text{O(g)} \rightarrow \text{MgO(s)} + \text{H}_2\text{(g)}\)

• Calcium, Strontium, and Barium with Cold Water: React with increasing vigour. Calcium fizzes steadily, the water turns cloudy as sparingly soluble \(\text{Ca(OH)}_2\) forms, and heat is released:
\(\text{Ca(s)} + 2\text{H}_2\text{O(l)} \rightarrow \text{Ca(OH)}_2\text{(aq/s)} + \text{H}_2\text{(g)}\)
Barium reacts vigorously, dissolving completely to give a clear, strongly alkaline solution of \(\text{Ba(OH)}_2\).

Common Mistake to Avoid: In exam questions asking for the reaction between magnesium and steam, students often incorrectly write \(\text{Mg(OH)}_2\) as the product. Remember: steam produces the oxide (\(\text{MgO}\)), whereas cold water produces the hydroxide (\(\text{Mg(OH)}_2\)).

3. Solubility Trends of Group 2 Compounds

One of the most heavily tested areas in CCEA AS Chemistry is the contrasting solubility patterns of Group 2 hydroxides versus Group 2 sulfates.

a) Solubility of Group 2 Hydroxides, \(\text{M(OH)}_2\)

Trend: Solubility increases down the group.

• \(\text{Mg(OH)}_2\) (Magnesium hydroxide): Sparingly soluble / Insoluble (Forms a white suspension in water, known as "Milk of Magnesia").
• \(\text{Ca(OH)}_2\) (Calcium hydroxide): Slightly soluble (A saturated aqueous solution is commonly called "limewater").
• \(\text{Sr(OH)}_2\) (Strontium hydroxide): Soluble.
• \(\text{Ba(OH)}_2\) (Barium hydroxide): Very soluble (Forms a strongly alkaline solution with high \(\text{OH}^-\) concentration).

b) Solubility of Group 2 Sulfates, \(\text{MSO}_4\)

Trend: Solubility decreases down the group.

• \(\text{MgSO}_4\) (Magnesium sulfate): Very soluble (Sold as Epsom salts).
• \(\text{CaSO}_4\) (Calcium sulfate): Slightly soluble.
• \(\text{SrSO}_4\) (Strontium sulfate): Insoluble.
• \(\text{BaSO}_4\) (Barium sulfate): Completely insoluble (Forms an instant, dense white precipitate).

Memory Aid for Solubilities

Use this simple mnemonic to keep the trends straight:
• Hydroxides = Higher solubility as you go down.
• Sulfates = Sinks (decreases) as you go down.

4. Thermal Stability of Group 2 Carbonates and Nitrates

When heated, solid Group 2 carbonates and nitrates undergo thermal decomposition (breaking down into simpler substances upon heating).

The Explanation: Cation Size and Polarising Power

Trend: Thermal stability increases down Group 2. This means you must heat the compound to a much higher temperature before it decomposes.

Why does this happen? Step-by-step:
1. All Group 2 cations carry the same \(2+\) charge.
2. Going down the group, the ionic radius of the cation increases (\(\text{Mg}^{2+} < \text{Ca}^{2+} < \text{Sr}^{2+} < \text{Ba}^{2+}\)).
3. Smaller cations (like \(\text{Mg}^{2+}\)) have a very high charge density.
4. High charge density means the cation has high polarising power: it strongly attracts the electron cloud of the nearby large, negative carbonate (\(\text{CO}_3^{2-}\)) or nitrate (\(\text{NO}_3^-\)) ion.
5. This polarises (distorts) the anion's covalent bonds (such as the \(\text{C–O}\) bond in carbonate), weakening it.
6. As a result, less heat energy is needed to break the weakened bond in \(\text{MgCO}_3\) compared to \(\text{BaCO}_3\), where the large \(\text{Ba}^{2+}\) ion has low charge density and causes very little distortion.

a) Thermal Decomposition of Carbonates

Group 2 carbonates decompose on heating to form the solid metal oxide and carbon dioxide gas:

\(\text{MCO}_3\text{(s)} \rightarrow \text{MO(s)} + \text{CO}_2\text{(g)}\)

Example: \(\text{CaCO}_3\text{(s)} \xrightarrow{\text{heat}} \text{CaO(s)} + \text{CO}_2\text{(g)}\)
• Observation: No colour change (both \(\text{CaCO}_3\) and \(\text{CaO}\) are white solids). The gas produced turns limewater cloudy.

b) Thermal Decomposition of Nitrates

Group 2 nitrates decompose on heating to form the metal oxide, toxic brown nitrogen dioxide gas (\(\text{NO}_2\)), and oxygen gas (\(\text{O}_2\)):

\(2\text{M(NO}_3)_2\text{(s)} \rightarrow 2\text{MO(s)} + 4\text{NO}_2\text{(g)} + \text{O}_2\text{(g)}\)

Example: \(2\text{Mg(NO}_3)_2\text{(s)} \rightarrow 2\text{MgO(s)} + 4\text{NO}_2\text{(g)} + \text{O}_2\text{(g)}\)
• Observations: White solid remains in the test tube, pungent brown fumes of \(\text{NO}_2\) gas are evolved, and a glowing splint will relight due to the production of \(\text{O}_2\).

Key Takeaway: Smaller cations = higher charge density = greater polarising power = more distortion of anion = lower thermal stability.

5. Flame Tests and Chemical Identification

a) Flame Tests for Group 2 Cations

When heated in a non-luminous (blue) Bunsen flame, outer electrons are excited to higher energy levels. When they fall back to ground state, they release energy in the form of visible light of characteristic wavelengths:

• Calcium (\(\text{Ca}^{2+}\)): Brick-red (or orange-red)
• Strontium (\(\text{Sr}^{2+}\)): Crimson-red
• Barium (\(\text{Ba}^{2+}\)): Apple-green

Did you know? Magnesium (\(\text{Mg}^{2+}\)) gives NO flame colour! The energy emitted when excited electrons in \(\text{Mg}^{2+}\) return to lower levels falls in the ultraviolet region of the electromagnetic spectrum, which is invisible to the human eye.

Standard Flame Test Procedure:
1. Dip a clean nichrome or platinum wire into concentrated hydrochloric acid (\(\text{HCl}\)) to convert metal compounds into volatile metal chlorides.
2. Hold the wire in a hot, blue Bunsen flame to ensure it is clean (no residual colour).
3. Dip the wire back into concentrated \(\text{HCl}\), touch a small sample of the solid salt, and place it into the edge of the blue Bunsen flame.
4. Observe and record the characteristic flame colour.

b) Testing for Sulfate Ions (\(\text{SO}_4^{2-}\))

Because barium sulfate (\(\text{BaSO}_4\)) is completely insoluble, this property is used as a standard qualitative test for sulfate ions in solution:

1. Add dilute hydrochloric acid (\(\text{HCl}\)) or dilute nitric acid (\(\text{HNO}_3\)) to the unknown solution. (Why? To remove any carbonate or sulfite impurities that would otherwise form a false white precipitate).
2. Add aqueous barium chloride solution, \(\text{BaCl}_2\text{(aq)}\) (or barium nitrate, \(\text{Ba(NO}_3)_2\text{(aq)}\)).
3. A thick white precipitate confirms the presence of sulfate ions:

\(\text{Ba}^{2+}\text{(aq)} + \text{SO}_4^{2-}\text{(aq)} \rightarrow \text{BaSO}_4\text{(s)}\)

6. Real-World Applications of Group 2 Compounds

You are expected to know several practical and industrial uses of Group 2 compounds:

• Magnesium Hydroxide, \(\text{Mg(OH)}_2\): Used as an antacid (Milk of Magnesia) to neutralize excess stomach hydrochloric acid (\(\text{HCl}\)) and relieve indigestion. Because it is sparingly soluble, it does not create an overly alkaline solution that would damage intestinal tissue:
\(\text{Mg(OH)}_2\text{(s)} + 2\text{HCl(aq)} \rightarrow \text{MgCl}_2\text{(aq)} + 2\text{H}_2\text{O(l)}\)

• Calcium Hydroxide, \(\text{Ca(OH)}_2\) (Slaked Lime): Used in agriculture to neutralise acidic soil and improve crop yield.

• Barium Sulfate, \(\text{BaSO}_4\) ("Barium Meal"): Used as a radiocontrast agent for X-ray imaging of the intestines. Free barium ions (\(\text{Ba}^{2+}\)) are highly toxic to humans, but \(\text{BaSO}_4\) is completely safe to ingest because its extreme insolubility prevents it from being absorbed into the bloodstream.

• Calcium Oxide (\(\text{CaO}\)) and Calcium Carbonate (\(\text{CaCO}_3\)): Used in power stations for flue gas desulfurisation. They react with acidic sulfur dioxide (\(\text{SO}_2\)) gas produced from burning coal, preventing it from entering the atmosphere and causing acid rain:
\(\text{CaO(s)} + \text{SO}_2\text{(g)} \rightarrow \text{CaSO}_3\text{(s)}\)

• Extraction of Titanium using Magnesium: Magnesium acts as a reducing agent to extract titanium from titanium(IV) chloride at high temperatures under an inert argon atmosphere:
\(\text{TiCl}_4\text{(g)} + 2\text{Mg(l)} \rightarrow \text{Ti(s)} + 2\text{MgCl}_2\text{(l)}\)

Quick Summary & Revision Checklist

Before sitting your exam, make sure you can answer the following:

• Can you write balanced equations with state symbols for \(\text{Mg}\) reacting with cold water vs. steam?
• Can you state the solubility trend for hydroxides (increases down) and sulfates (decreases down)?
• Can you explain why carbonates become more thermally stable down the group using the terms ionic radius, charge density, and polarising power?
• Do you know your flame colours? (\(\text{Ca}^{2+}\) = brick-red, \(\text{Sr}^{2+}\) = crimson, \(\text{Ba}^{2+}\) = apple-green, \(\text{Mg}^{2+}\) = no colour).
• Can you state the ionic equation for the sulfate test with \(\text{Ba}^{2+}\)?

Quick check

Can you answer these now?

Open each question to check the key ideas from this chapter.

What is the general trend in first and second ionisation energies down Group 2, and why?

Ionisation energies decrease down the group because atomic radius and electron shielding increase, which outweighs the increase in nuclear charge and weakens the attraction on outer electrons.

Why do melting points generally decrease down Group 2?

Down the group, the ionic radius of the \(\text{M}^{2+}\) cation increases while the charge stays constant, decreasing the charge density and weakening the metallic bonding with delocalised electrons.

What are the different products when magnesium reacts with cold water versus steam?

With cold water, magnesium reacts slowly to form magnesium hydroxide (\(\text{Mg(OH)}_2\)) and \(\text{H}_2\); with steam, it reacts vigorously to form magnesium oxide (\(\text{MgO}\)) and \(\text{H}_2\).

What are the solubility trends for Group 2 hydroxides and sulfates down the group?

Hydroxide solubility increases down the group, whereas sulfate solubility decreases down the group.

Why does the thermal stability of Group 2 carbonates increase down the group?

Larger cations down the group have lower charge density and lower polarising power, causing less distortion of the carbonate ion's \(\text{C–O}\) bonds.

What are the products and observations during the thermal decomposition of Group 2 nitrates?

They decompose into a metal oxide, brown nitrogen dioxide gas (\(\text{NO}_2\)), and oxygen gas (\(\text{O}_2\)).

What flame colours are produced by \(\text{Ca}^{2+}\), \(\text{Sr}^{2+}\), \(\text{Ba}^{2+}\), and \(\text{Mg}^{2+}\)?

\(\text{Ca}^{2+}\) gives brick-red, \(\text{Sr}^{2+}\) gives crimson-red, \(\text{Ba}^{2+}\) gives apple-green, and \(\text{Mg}^{2+}\) gives no flame colour.

How do you test for sulfate ions (\(\text{SO}_4^{2-}\)) in solution?

Acidify the sample with dilute \(\text{HCl}\) (or \(\text{HNO}_3\)), then add aqueous \(\text{BaCl}_2\); a white precipitate of \(\text{BaSO}_4\) confirms sulfate ions.

Why is barium sulfate (\(\text{BaSO}_4\)) safe to ingest for medical X-ray imaging despite \(\text{Ba}^{2+}\) ions being toxic?

\(\text{BaSO}_4\) is completely insoluble in water, preventing toxic \(\text{Ba}^{2+}\) ions from being absorbed into the bloodstream.

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