Group 2 chemistry

Q: What is the general valence electron configuration of Group 2 elements and what ion do they typically form?

Group 2 elements have an outer electron configuration of s^2 and lose these two electrons to form stable \(2+\) cations (\(\text{M}^{2+}\)).

Q: Why do the first and second ionisation energies decrease down Group 2?

Atomic radius and electron shielding increase down the group, which outweighs the increase in nuclear charge and weakens the attraction between the nucleus and outer electrons.

Q: Why does the chemical reactivity of Group 2 elements increase down the group?

Reactivity increases because the sum of the first two ionisation energies decreases down the group, making it easier for larger atoms to lose their valence electrons.

Q: What is the difference between the reaction of magnesium with cold water versus its reaction with steam?

Magnesium reacts very slowly with cold water to form magnesium hydroxide, \(\text{Mg(OH)}_2\), but burns vigorously in steam to form magnesium oxide, \(\text{MgO}\), and hydrogen gas.

Q: How does the solubility of Group 2 hydroxides vary down the group, and how does this affect solution pH?

The solubility of Group 2 hydroxides increases down the group, which results in a higher concentration of \(\text{OH}^-\text{(aq)}\) ions and a higher \(\text{pH}\) for saturated solutions.

Q: How is aqueous barium chloride used to test for sulfate ions, and what is the role of hydrochloric acid in the test?

Hydrochloric acid removes interfering carbonate or sulfite ions, after which adding \(\text{BaCl}_2\text{(aq)}\) precipitates sulfate ions as insoluble white barium sulfate, \(\text{BaSO}_4\text{(s)}\).

Q: Why can toxic barium sulfate safely be administered internally to patients as a "barium meal" for X-ray imaging?

\(\text{BaSO}_4\) is completely insoluble in water, preventing toxic \(\text{Ba}^{2+}\) ions from entering the bloodstream while remaining opaque to X-rays.

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

Cations become larger down the group, reducing their charge density and polarising power, which causes less distortion and weakening of the \(\text{C}-\text{O}\) bonds in the carbonate ion.

Q: What characteristic flame colours are produced by calcium, strontium, and barium cations?

\(\text{Ca}^{2+}\) gives a brick-red flame, \(\text{Sr}^{2+}\) gives a red (crimson) flame, and \(\text{Ba}^{2+}\) gives an apple-green flame.

Chemistry B (Salters) - H433 · Cambridge OCR A Level · 12 min read · Updated: Oct 22, 2026

Introduction to Group 2: The Alkaline Earth Metals

Welcome to Group 2 chemistry! In your study of the periodic table within the Elements of life storyline, Group 2 elements play a starring role. You have likely encountered them in limestone cliffs, indigestion tablets, construction cement, and even bright fireworks displays. These elements are known as the alkaline earth metals, and they occupy the s-block of the periodic table.

Don't worry if periodic trends felt overwhelming in earlier topics. Group 2 is one of the most logical and predictable families in the entire periodic table. Once you understand the core patterns — how atomic radius, shielding, and nuclear attraction change down the group — every reaction and trend falls neatly into place.

Note: For broader periodic patterns across periods and general electronic principles, refer to the companion chapters on "Periodicity and the periodic table" and "Atomic structure and electron configurations".

Electronic Structure and the Formation of \(2+\) Ions

Every element in Group 2 has two electrons in its outermost s sub-shell, giving them the general outer valence electron configuration of \(\text{n}s^2\):

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

In all their characteristic chemical reactions, Group 2 elements act as reducing agents. They lose their two outer \(s\) electrons to form stable, full-shell cations with a \(2+\) charge:

\(\text{M} \rightarrow \text{M}^{2+} + 2\text{e}^-\)

Because they give away electrons, they are oxidised (their oxidation state increases from \(0\) to \(+2\)), while reducing another species.

Trends in Physical Properties Down Group 2

1. Atomic Radius Increases Down the Group
As you move from beryllium down to barium, each successive element possesses an extra principal quantum energy level (electron shell). This increases the physical distance between the nucleus and the outer valence electrons, so the atomic radius increases.

2. First and Second Ionisation Energies Decrease Down the Group
Recall that ionisation energy is the energy needed to remove an electron from each atom or ion in one mole of gaseous atoms or ions. When a Group 2 metal reacts, both the first and second ionisation energies must be supplied:

First ionisation energy: \(\text{M(g)} \rightarrow \text{M}^+\text{(g)} + \text{e}^-\)
Second ionisation energy: \(\text{M}^+\text{(g)} \rightarrow \text{M}^{2+}\text{(g)} + \text{e}^-\)

As you descend the group:
• More electron shielding: There are more filled inner shells between the positive nucleus and the outermost electrons.
• Greater atomic radius: Outer electrons are further away from the nucleus.
• Although the nuclear charge (number of protons) increases, the extra shielding and distance far outweigh this factor.
• Consequently, the electrostatic attraction between the positive nucleus and the outermost electrons becomes progressively weaker, meaning less energy is required to remove the outer two electrons.

Key Takeaway: Down Group 2, outer electrons are held less tightly by the nucleus because of increased distance and shielding. As a result, ionisation energies decrease, which directly drives the trend in reactivity.

Reactivity of Group 2 Elements

Since Group 2 metals react by losing their two valence electrons, the easier it is to remove those electrons, the faster and more vigorously the metal reacts. Because the sum of the first two ionisation energies decreases down the group, chemical reactivity increases down Group 2 from magnesium to barium.

1. Reactions with Oxygen

All Group 2 metals react with oxygen gas to form white, ionic solid metal oxides where the metal has an oxidation state of \(+2\) and oxygen is \(-2\):

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

For example, when magnesium ribbon is heated in air, it burns with a characteristic, brilliant white flame to leave behind white magnesium oxide powder:

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

Barium reacts even more vigorously, igniting readily upon heating and reacting slowly with oxygen even at room temperature.

2. Reactions with Water

Group 2 metals react with liquid water to produce an alkaline metal hydroxide and hydrogen gas:

\(\text{M(s)} + 2\text{H}_2\text{O(l)} \rightarrow \text{M(OH)}_2\text{(aq or s)} + \text{H}_2\text{(g)}\)

Notice the change in oxidation states: the metal is oxidised from \(0\) to \(+2\), and hydrogen in water is reduced from \(+1\) to \(0\). The vigor of this reaction increases visibly down the group:

• Magnesium (\(\text{Mg}\)): Reacts exceptionally slowly with cold liquid water. Only a few tiny bubbles of hydrogen gas appear over several days because a thin, insoluble layer of magnesium hydroxide forms on the metal surface, stopping further contact.
• Calcium (\(\text{Ca}\)): Reacts steadily with cold water. The liquid fizzes vigorously with bubbles of hydrogen gas, the test tube warms up (exothermic), and a cloudy white suspension of calcium hydroxide forms:
\(\text{Ca(s)} + 2\text{H}_2\text{O(l)} \rightarrow \text{Ca(OH)}_2\text{(s or aq)} + \text{H}_2\text{(g)}\)
• Strontium (\(\text{Sr}\)) and Barium (\(\text{Ba}\)): React with increasing speed and violence with cold water, dissolving rapidly to give clear, strongly alkaline solutions as their hydroxides are significantly more soluble.

Special Case: Magnesium with Steam

A favourite exam question tests the stark contrast between magnesium reacting with cold water versus magnesium reacting with hot steam. When heated magnesium ribbon is exposed to steam (water vapour), it burns brilliantly with a bright white light, reacting rapidly:

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

Crucial Distinction to Remember:
• Reaction with cold liquid water gives the hydroxide: \(\text{Mg(OH)}_2\)
• Reaction with steam provides enough thermal energy to decompose any hydroxide, forming the white solid oxide: \(\text{MgO}\)

3. Reactions of Group 2 Oxides and Hydroxides with Acids

Group 2 oxides and hydroxides are basic substances. They neutralise acids to produce a salt and water (acting as Bronsted-Lowry proton acceptors):

Oxide with acid:
\(\text{MO(s)} + 2\text{HCl(aq)} \rightarrow \text{MCl}_2\text{(aq)} + \text{H}_2\text{O(l)}\)
For example: \(\text{CaO(s)} + 2\text{HNO}_3\text{(aq)} \rightarrow \text{Ca(NO}_3)_2\text{(aq)} + \text{H}_2\text{O(l)}\)

Hydroxide with acid:
\(\text{M(OH)}_2\text{(aq or s)} + 2\text{HCl(aq)} \rightarrow \text{MCl}_2\text{(aq)} + 2\text{H}_2\text{O(l)}\)
Ionic equation: \(\text{H}^+\text{(aq)} + \text{OH}^-\text{(aq)} \rightarrow \text{H}_2\text{O(l)}\)

Everyday Connection: Indigestion remedies often contain magnesium hydroxide, \(\text{Mg(OH)}_2\) (known as "milk of magnesia"), or calcium carbonate, \(\text{CaCO}_3\). Because \(\text{Mg(OH)}_2\) is sparingly soluble, it produces a low, gentle concentration of hydroxide ions that neutralises excess hydrochloric acid in the stomach without causing caustic burns to delicate tissues.

Key Takeaway: Reactivity with water increases down Group 2 as outer electrons become easier to remove. Magnesium reacts very slowly with cold water to form \(\text{Mg(OH)}_2\), but burns brightly in steam to form \(\text{MgO}\).

Solubility Trends of Group 2 Compounds

Understanding solubility trends is crucial for both theoretical exam questions and practical qualitative analysis. Students often mix up these two opposite trends, so study them carefully!

1. Solubility of Group 2 Hydroxides: Increases Down the Group

The solubility of Group 2 hydroxides in water increases down the group:

• \(\text{Mg(OH)}_2\) is sparingly soluble (almost insoluble, forming a suspension). A saturated solution has a mild \(\text{pH}\) of around \(9\) to \(10\).
• \(\text{Ca(OH)}_2\) is slightly soluble. A saturated aqueous solution of calcium hydroxide is known as limewater, used to test for carbon dioxide gas (turning milky as insoluble \(\text{CaCO}_3\) precipitates). Its \(\text{pH}\) is around \(11\).
• \(\text{Sr(OH)}_2\) is moderately soluble.
• \(\text{Ba(OH)}_2\) is readily soluble, dissolving to give a high concentration of free \(\text{OH}^-\text{(aq)}\) ions and forming a strongly alkaline solution with a \(\text{pH}\) around \(13\).

Because solubility increases down the group, the concentration of dissolved hydroxide ions in saturated solutions increases, meaning the \(\text{pH}\) of saturated solutions increases down the group.

2. Solubility of Group 2 Sulfates: Decreases Down the Group

The solubility of Group 2 sulfates in water decreases down the group:

• \(\text{MgSO}_4\) is freely soluble (used as Epsom salts).
• \(\text{CaSO}_4\) is slightly soluble.
• \(\text{SrSO}_4\) is insoluble.
• \(\text{BaSO}_4\) is completely insoluble in water.

Memory Trick:
Keep the initials in mind:
• Hydroxide → Higher solubility as you go down (both start with H).
• Sulfate → Sinks (decreases) as you go down (both start with S).

Practical Importance: Testing for Sulfate Ions and Medicine

The insolubility of barium sulfate provides an essential laboratory test for sulfate ions, \(\text{SO}_4^{2-}\):

1. Take the unknown solution and add dilute hydrochloric acid, \(\text{HCl(aq)}\). This acidifies the sample to destroy and eliminate any carbonate ions (\(\text{CO}_3^{2-}\)) or sulfite ions that would otherwise form a misleading white precipitate with barium.
2. Add aqueous barium chloride, \(\text{BaCl}_2\text{(aq)}\) (or barium nitrate, \(\text{Ba(NO}_3)_2\)).
3. A dense white precipitate of barium sulfate confirms the presence of sulfate ions:

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

Barium Meals in Medicine:
Although soluble barium compounds are highly toxic to humans, patients undergoing gastrointestinal X-ray imaging can safely swallow a suspension of \(\text{BaSO}_4\) (a "barium meal"). Why? Barium sulfate is so utterly insoluble in water that virtually no toxic \(\text{Ba}^{2+}\) ions are absorbed into the bloodstream. Furthermore, barium's heavy nucleus absorbs X-rays strongly, rendering the soft digestive tract clearly visible on X-ray plates.

Key Takeaway: Hydroxides become more soluble down Group 2 (increasing \(\text{pH}\)); sulfates become less soluble down Group 2, making \(\text{BaSO}_4\) an insoluble white solid used for sulfate testing and medical imaging.

Thermal Stability of Group 2 Carbonates

When solid Group 2 carbonates are heated strongly, they do not melt; instead, they undergo thermal decomposition to produce a solid metal oxide and carbon dioxide gas:

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

For example, the industrial decomposition of limestone (calcium carbonate):

\(\text{CaCO}_3\text{(s)} \rightarrow \text{CaO(s)} + \text{CO}_2\text{(g)}\)

Trend in Thermal Stability

The thermal stability of Group 2 carbonates increases down the group. Barium carbonate requires much more intense and prolonged heating to decompose than magnesium carbonate:

• \(\text{MgCO}_3\): decomposes readily around \(300^\circ\text{C}\)
• \(\text{CaCO}_3\): decomposes around \(900^\circ\text{C}\)
• \(\text{SrCO}_3\): decomposes around \(1200^\circ\text{C}\)
• \(\text{BaCO}_3\): decomposes only above \(1360^\circ\text{C}\)

Explaining the Trend: Charge Density and Polarisation

This is a classic extended response question in OCR Chemistry B (Salters). Follow this logical step-by-step reasoning:

Step 1: Size and charge of the cation
All Group 2 metal cations carry the exact same charge (\(2+\)). However, as you descend the group from \(\text{Mg}^{2+}\) to \(\text{Ba}^{2+}\), the ionic radius increases significantly because of additional electron shells.

Step 2: Charge density decreases
Because the same \(+2\) charge is spread over a progressively larger ionic volume, the charge density of the cation decreases down the group. The smaller \(\text{Mg}^{2+}\) ion has a very high charge density, whereas the larger \(\text{Ba}^{2+}\) ion has a much lower charge density.

Step 3: Polarising power decreases
A cation with high charge density (\(\text{Mg}^{2+}\)) exerts a strong electrostatic pull on the electron cloud of the large, negative carbonate anion (\(\text{CO}_3^{2-}\)). This distorts the electron cloud of the carbonate ion — an effect called polarisation.

Step 4: Weakening of the carbon-oxygen bond
The polarisation of the carbonate ion pulls electron density away from one of the \(\text{C}-\text{O}\) bonds towards the metal cation. This weakens that specific \(\text{C}-\text{O}\) covalent bond within the carbonate ion, making it easy to break off a \(\text{CO}_2\) molecule when heated.

Conclusion:
Down the group, the larger cations have lower charge density, cause less polarisation of the carbonate electron cloud, and do not weaken the \(\text{C}-\text{O}\) bond as much. Therefore, more thermal energy is required to break the bonds and decompose the compound, so thermal stability increases down the group.

Note: Group 2 nitrates show the exact same trend: thermal stability increases down the group for the identical polarising power reason.

Key Takeaway: Down Group 2, cations get larger, their charge density falls, and their polarising power diminishes. As a result, the carbonate anion is less distorted, and more thermal energy is needed for decomposition.

Flame Tests for Group 2 Cations

In qualitative analysis, flame tests provide a fast, convenient method to identify certain metal cations based on characteristic emission colours met during your study of atomic spectra.

How a Flame Test is Performed (Section 1.2 / Practical Skills)

1. Dip a clean nichrome (or platinum) wire loop into concentrated hydrochloric acid (\(\text{HCl(aq)}\)).
2. Hold the wire in a hot, non-luminous (roaring blue) Bunsen flame until no colour is imparted, confirming the wire is clean.
3. Dip the clean wire back into concentrated hydrochloric acid so that a tiny sample of the solid metal salt sticks to it. (Using \(\text{HCl}\) converts metal salts to volatile metal chlorides, which vaporise easily into the flame).
4. Place the wire into the edge of the hot blue flame and observe the characteristic colour produced.

Characteristic Flame Colours

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

Why doesn't Magnesium give a distinctive flame colour?
Magnesium compounds do not produce a distinctive visible flame colour in a standard flame test. The energy supplied by the Bunsen burner excites electrons to higher energy states; when they fall back down, the emission wavelengths emitted fall outside the visible spectrum (chiefly in the ultraviolet region), so no distinct visible color is observed by the human eye.

Common Exam Mistake: Do not confuse the flame colour of burning magnesium metal with a flame test of a magnesium salt! Burning magnesium metal gives an intense, blinding white light due to the high temperature of combustion. But in a qualitative flame test of a magnesium compound (such as \(\text{MgCl}_2\)), there is no distinctive flame colour.

Key Takeaway: Flame tests identify cations: \(\text{Ca}^{2+}\) is brick-red, \(\text{Sr}^{2+}\) is red/crimson, and \(\text{Ba}^{2+}\) is apple-green.

Summary and Quick Revision Guide

Here is your at-a-glance guide to mastering Group 2 chemistry questions:

1. Reactivity with water and oxygen:
Increases down the group (\(\text{Mg} < \text{Ca} < \text{Sr} < \text{Ba}\)). Driven by lower ionisation energies due to larger atomic radii and greater electron shielding.

2. Solubility of hydroxides, \(\text{M(OH)}_2\):
Increases down the group (\(\text{Mg(OH)}_2\) is sparingly soluble; \(\text{Ba(OH)}_2\) is readily soluble). Resulting saturated solutions have higher \(\text{pH}\) down the group.

3. Solubility of sulfates, \(\text{MSO}_4\):
Decreases down the group (\(\text{MgSO}_4\) is soluble; \(\text{BaSO}_4\) is insoluble). Insoluble \(\text{BaSO}_4\) forms the basis of the test for sulfate ions.

4. Thermal stability of carbonates, \(\text{MCO}_3\):
Increases down the group. Cation radius increases → charge density falls → less polarisation of the \(\text{CO}_3^{2-}\) electron cloud → \(\text{C}-\text{O}\) bond weakened less → more heat needed to decompose.

5. Flame colours:
\(\text{Ca}^{2+}\) = brick-red; \(\text{Sr}^{2+}\) = red/crimson; \(\text{Ba}^{2+}\) = apple-green.

Quick check

Can you answer these now?

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

What is the general valence electron configuration of Group 2 elements and what ion do they typically form?

Group 2 elements have an outer electron configuration of s^2 and lose these two electrons to form stable \(2+\) cations (\(\text{M}^{2+}\)).

Why do the first and second ionisation energies decrease down Group 2?

Atomic radius and electron shielding increase down the group, which outweighs the increase in nuclear charge and weakens the attraction between the nucleus and outer electrons.

Why does the chemical reactivity of Group 2 elements increase down the group?

Reactivity increases because the sum of the first two ionisation energies decreases down the group, making it easier for larger atoms to lose their valence electrons.

What is the difference between the reaction of magnesium with cold water versus its reaction with steam?

Magnesium reacts very slowly with cold water to form magnesium hydroxide, \(\text{Mg(OH)}_2\), but burns vigorously in steam to form magnesium oxide, \(\text{MgO}\), and hydrogen gas.

How does the solubility of Group 2 hydroxides vary down the group, and how does this affect solution pH?

The solubility of Group 2 hydroxides increases down the group, which results in a higher concentration of \(\text{OH}^-\text{(aq)}\) ions and a higher \(\text{pH}\) for saturated solutions.

How is aqueous barium chloride used to test for sulfate ions, and what is the role of hydrochloric acid in the test?

Hydrochloric acid removes interfering carbonate or sulfite ions, after which adding \(\text{BaCl}_2\text{(aq)}\) precipitates sulfate ions as insoluble white barium sulfate, \(\text{BaSO}_4\text{(s)}\).

Why can toxic barium sulfate safely be administered internally to patients as a "barium meal" for X-ray imaging?

\(\text{BaSO}_4\) is completely insoluble in water, preventing toxic \(\text{Ba}^{2+}\) ions from entering the bloodstream while remaining opaque to X-rays.

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

Cations become larger down the group, reducing their charge density and polarising power, which causes less distortion and weakening of the \(\text{C}-\text{O}\) bonds in the carbonate ion.

What characteristic flame colours are produced by calcium, strontium, and barium cations?

\(\text{Ca}^{2+}\) gives a brick-red flame, \(\text{Sr}^{2+}\) gives a red (crimson) flame, and \(\text{Ba}^{2+}\) gives an apple-green flame.

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