Chapter Summary

• Group 1 (alkali metals) have 1 outer electron; Group 2 (alkaline earth metals) have 2 outer electrons. Both lose these electrons to form ions with noble gas configuration, acting as reducing agents.
• Going down either group: ionisation energy decreases → reactivity increases (bigger atoms, more shielding, weaker nuclear pull on the outer electron).
• Group 1 and 2 metals react with oxygen (→ oxide), chlorine (→ chloride), and water (→ hydroxide + hydrogen) — reactions get more vigorous down the group.
• Metal oxides + water → hydroxides (alkaline solutions). Metal oxides/hydroxides + acid → salt + water (neutralisation).
• Group 2 hydroxides get more soluble down the group (more alkaline solutions); Group 2 sulfates get less soluble down the group.
• Thermal stability of carbonates and nitrates increases down the group — smaller, more charge-dense cations polarise (distort) the big anion more, weakening it so it decomposes more easily at lower temperatures.
• Flame tests identify Group 1/2 metal ions by characteristic flame colours (e.g. lilac = K⁺, brick red = Ca²⁺).
• Qualitative test-tube tests exist for ammonium (NH₄⁺), carbonate/hydrogencarbonate (CO₃²⁻/HCO₃⁻), and sulfate (SO₄²⁻) ions.

1. Ionisation Energy — Groups 1 & 2

Think of an atom's outer electron like a satellite orbiting a planet. The "pull" that keeps it in orbit is the attraction between the negative electron and the positive nucleus. Ionisation energy is just a measure of how hard you'd have to yank to remove that electron completely.

Every Group 1 atom has exactly 1 electron in its outermost shell, and every Group 2 atom has exactly 2 electrons in its outermost shell. When these metals react, they don't want to gain electrons to fill the shell (that would take way too much energy) — instead, it's much easier to lose those 1 or 2 outer electrons entirely. Once they're gone, the shell underneath — which is already completely full — becomes the new outer shell. That's a stable, noble-gas-like electron arrangement, and it's the "reward" that makes losing electrons energetically worthwhile.

📉Why ionisation energy falls going down the group

Here's the part students often get backwards: nuclear charge increases going down the group (more protons = more positive charge in the nucleus), so you might expect the electron to be held on tighter. But three things overpower that increased pull:

Because the distance and shielding effects win out over the extra nuclear charge, less energy is needed to remove the outer electron(s) as you go down the group. This is the root cause of almost everything else in this chapter — reactivity trends, reaction vigour, all of it traces back to this one idea.

2. Reactivity Trends & Reactions

⚡Why reactivity increases down the group

This follows directly from Section 1. If it takes less energy to remove the outer electron(s), the metal will react faster and more vigorously whenever a reaction requires it to lose those electrons — which is every single reaction in this chapter. So: lower ionisation energy → more reactive metal, full stop.

🔥Reaction with Oxygen

Group 1 metals tarnish (go dull) in air because they react with oxygen to form a metal oxide coating on the surface. This tarnishing happens faster down the group.

For Group 2: 2M(s) + O₂(g) → 2MO(s) where M is any Group 2 metal. Note that Sr and Ba can also form a peroxide, MO₂.

🧪Reaction with Chlorine

Group 1 metals burn in chlorine gas the same way they burn in oxygen — sodium, for example, burns with an intense orange flame in both gases. In each case a white solid (the simple metal chloride) is formed.

Group 2 metals behave the same way, e.g. Mg(s) + Cl₂(g) → MgCl₂(s).

💧Reaction with Water

This is where the reactivity trend is most dramatic and most commonly tested.

For Group 2, note the sharp contrast in speed: Mg reacts extremely slowly with cold water (barely visible), while Ca reacts fairly vigorously in a clearly exothermic reaction. This is the same reactivity trend at work — Mg is higher up the group so has a higher ionisation energy and is less reactive than Ca below it.

Interesting twist: magnesium barely reacts with cold water because the Mg(OH)₂ that forms is only slightly soluble and coats the surface, slowing the reaction. But heat magnesium in steam instead, and it reacts vigorously:

Notice the product is different — MgO (the oxide) forms with steam, not Mg(OH)₂ (the hydroxide) which forms with cold liquid water.

3. Oxides & Hydroxides

There's a neat two-step chain reaction pattern running through this whole topic:

Metal oxides are basic (except BeO, which is amphoteric — meaning it can react as both an acid and a base). When metal oxides dissolve in water, they release hydroxide ions, making the solution alkaline:

Calcium oxide reacting with water is a classic example, and its product has a special name you must know:

🧫Neutralisation Reactions

Both metal oxides and metal hydroxides neutralise dilute acids to form a salt + water — this is standard acid-base chemistry:

Worked examples:

📊Group 2 Oxides Reacting with Water

4. Group 2 Hydroxides & Sulfates: Solubility Trends

These two trends run in opposite directions — a classic exam trap if you mix them up. Get the pattern locked in with this memory device:

Why does this matter for alkalinity? When a hydroxide dissolves, it releases OH⁻ ions:

More soluble hydroxide → more OH⁻ ions released into solution → higher pH → more alkaline. This directly explains why the reaction of Group 2 oxides with water gets more vigorous and more alkaline as you go down the group (see the CaO/BaO table above).

5. Nitrates & Carbonates: Thermal Stability

Thermal decomposition = breaking a compound apart into two or more substances using heat. This section is about which carbonates and nitrates break down easily with heat, and which need much higher temperatures.

🪨Carbonates

Lithium is the odd one out in Group 1 — it decomposes on heating, while the rest of Group 1's carbonates (Na, K, Rb, Cs) don't decompose at normal Bunsen burner temperatures.

All Group 2 carbonates do decompose (unlike most of Group 1), releasing CO₂ and leaving the metal oxide behind:

Example: MgCO₃ —heat→ MgO + CO₂

💥Nitrates

Lithium nitrate is again the exception in Group 1 — it decomposes fully to the oxide, releasing toxic brown NO₂ gas and oxygen:

The rest of Group 1 (Na through Cs) decompose less completely — only as far as the metal nitrite (NO₂⁻), with no brown NO₂ gas produced:

All Group 2 nitrates decompose fully to the oxide, releasing both oxygen and toxic NO₂ gas (like lithium does):

🎯The Explanation: Ionic Polarisation

This is the concept examiners love to test, so let's build it up carefully using an analogy. Picture the carbonate or nitrate ion (CO₃²⁻ or NO₃⁻) as a soft balloon full of delocalised electrons floating evenly around it. Now imagine a small, strongly-charged positive ion sitting right next to this balloon — that positive charge will pull and distort the electron cloud towards itself, squashing the balloon's shape. This distorting effect is called polarisation.

So the rule is simple once you see the mechanism: smaller, higher-charge cations polarise the anion more strongly, weakening its internal bonds and making it decompose at a lower temperature. Bigger cations at the bottom of the group have their charge spread more thinly, so they barely distort the anion, and much higher temperatures are needed to break it apart.

🧪Testing for the Gases Produced

6. Flame Tests

Since all Group 1 (and most Group 2) compounds are soluble, we can't use precipitation to identify their cations. Instead, we exploit something unique to each metal ion: the specific colour of light it emits when heated in a flame.

🔥Method

1. Dip a loop of unreactive wire (nichrome or platinum) into concentrated acid, then hold it in the blue Bunsen flame until there's no colour change — this cleans the loop.
2. This cleaning step matters because the test only works with one type of ion present — mixing two or more ions makes the colours blend, giving an unreliable result.
3. Dip the clean loop into the solid sample, then place it at the edge of the blue Bunsen flame.
4. Avoid letting the wire get hot enough to glow red — this could be mistaken for a flame colour result.

🎨Flame Colours to Memorise

7. Qualitative Tests

Three test-tube reactions you need to know cold: testing for ammonium ions, carbonate/hydrogencarbonate ions, and sulfate ions. If the sample is a solid, dissolve it in deionised water first.

🧪Testing for Ammonium Ions (NH₄⁺)

Method: Add ~10 drops of the ammonium solution to a clean test tube, then add ~10 drops of sodium hydroxide.

Swirl to mix, then gently heat the test tube in a water bath (beaker of water over a Bunsen). Hold a piece of damp red litmus paper near the mouth of the tube with tongs.

Alternative confirmation: hold a rod dipped in concentrated HCl near the mouth of the tube — a white smoke of ammonium chloride will form:

🧫Testing for Carbonate & Hydrogencarbonate Ions

Method: Add ~1 cm³ dilute HCl to a test tube. Add an equal amount of the carbonate solution using a clean pipette, then immediately attach a bung with a delivery tube leading into a second test tube containing limewater.

For hydrogencarbonates, the equation is very similar but only needs 1 H⁺ per ion:

⚗️Testing for Sulfate Ions (SO₄²⁻)

Method: Acidify the sample with dilute HCl, then add a few drops of aqueous barium chloride.

What to Memorise

Concepts Checklist

Exam Tips & Common Mistakes