Welcome to Metals and the Reactivity Series
Have you ever wondered why gold stays shiny for thousands of years in ancient jewellery, while an iron nail left outside rusts in just a few days? Or why magnesium burns with a blinding white flare, while copper simply turns dull and black when heated?
It all comes down to chemical reactivity. In this chapter of Unit 2, we will explore how different metals behave, arrange them into a "league table" called the Reactivity Series, and learn how to write balanced symbol and ionic equations for their reactions with oxygen, water, steam, and salt solutions.
Don't worry if this seems like a lot to memorise at first! We will break down every reaction step-by-step with simple patterns, memory tricks, and clear examiner tips.
---1. The Reactivity Series & What Drives Reactivity
What is the Reactivity Series?
The reactivity series places metals in order of their chemical reactivity, from the most reactive at the top to the least reactive at the bottom.
Here is the official CCEA series you need to know:
\( \text{Potassium (K)} > \text{Sodium (Na)} > \text{Calcium (Ca)} > \text{Magnesium (Mg)} > \text{Aluminium (Al)} > \text{Zinc (Zn)} > \text{Iron (Fe)} > \text{Copper (Cu)} \)
Note: Silver (\(\text{Ag}\)) and Gold (\(\text{Au}\)) sit at the very unreactive bottom, while non-metals like Carbon (\(\text{C}\)) and Hydrogen (\(\text{H}\)) are useful comparison benchmarks.
Memory Mnemonic
To keep the order firmly in your head, use this handy sentence:
Please Send Charlie's Monkeys And Zebras In Cages
Potassium \(\rightarrow\) Sodium \(\rightarrow\) Calcium \(\rightarrow\) Magnesium \(\rightarrow\) Aluminium \(\rightarrow\) Zinc \(\rightarrow\) Iron \(\rightarrow\) Copper
Why Are Some Metals More Reactive Than Others?
Metals react by losing their outer-shell electrons to form positive ions (cations):
\( \text{M} \rightarrow \text{M}^{n+} + ne^- \)
The easier it is for a metal atom to lose its valence electron(s) (the ease of oxidation), the more reactive the metal is. Potassium loses its single outer electron extremely easily, making it violently reactive. Copper holds onto its outer electrons much more tightly, making it unreactive.
The Curious Case of Aluminium (\(\text{Al}\))
Did you know? Looking at its position, aluminium should react rapidly. However, in everyday life, aluminium foil and window frames do not react visibly with air or water. This is because aluminium quickly forms a tough, impervious, unreactive surface layer of aluminium oxide (\(\text{Al}_2\text{O}_3\)) that seals the metal underneath from further attack.
Section Key Takeaway: Reactivity measures how easily a metal atom loses electrons to become a positive ion. Potassium is at the top (easiest to oxidise); copper is near the bottom.
---2. Reactions of Metals with Air / Oxygen
When metals are heated in air, they undergo an oxidation reaction by reacting with oxygen gas (\(\text{O}_2\)) to form metal oxides.
General Word Equation:
\( \text{Metal} + \text{Oxygen} \rightarrow \text{Metal Oxide} \)
Key Observations and Equations
Examiners love asking for flame colours and product appearances. Here is what you will observe for each metal:
• Potassium (\(\text{K}\)): Burns vigorously with a distinctive lilac flame, forming a white solid (\(\text{K}_2\text{O}\)).
• Sodium (\(\text{Na}\)): Burns vigorously with a bright yellow/orange flame, forming a white solid (\(\text{Na}_2\text{O}\)).
• Calcium (\(\text{Ca}\)): Burns with a brick-red flame, producing a white solid (\(\text{CaO}\)).
• Magnesium (\(\text{Mg}\)): Burns with a brilliant/blinding white light, leaving behind a white powder / solid (\(\text{MgO}\)):
\( 2\text{Mg (s)} + \text{O}_2\text{ (g)} \rightarrow 2\text{MgO (s)} \)
• Aluminium (\(\text{Al}\)): Powdered aluminium burns with a bright white flame to form white \( \text{Al}_2\text{O}_3 \).
• Zinc (\(\text{Zn}\)): Burns moderately to form zinc oxide (\(\text{ZnO}\)), which has a unique property: it is yellow when hot and turns white on cooling.
• Iron (\(\text{Fe}\)): Iron filings glow with orange sparks, leaving a dark grey/black solid (\(\text{Fe}_3\text{O}_4\) or \( \text{Fe}_2\text{O}_3 \)).
• Copper (\(\text{Cu}\)): Does not burn. The surface tarnishes and becomes coated in a black layer of copper(II) oxide:
\( 2\text{Cu (s)} + \text{O}_2\text{ (g)} \rightarrow 2\text{CuO (s)} \)
Section Key Takeaway: More reactive metals burn vigorously with distinctive flame colours (Potassium = lilac, Sodium = yellow/orange, Calcium = brick-red, Magnesium = blinding white). Less reactive metals like copper do not burn, but form a black oxide coating.
---3. Reactions with Cold Water vs. Steam
A classic exam trap is confusing what forms when a metal reacts with cold liquid water versus when it reacts with hot gaseous steam. Let's compare them clearly.
Part A: Reactions with Cold Water
Only the most reactive metals react with cold water to form a metal hydroxide and hydrogen gas.
General Word Equation:
\( \text{Metal} + \text{Water (liquid)} \rightarrow \text{Metal Hydroxide} + \text{Hydrogen} \)
1. Potassium (\(\text{K}\)):
• Floats on the water surface.
• Melts into a small silvery ball.
• Fizzes vigorously as hydrogen gas is produced.
• Ignites spontaneously and burns with a lilac flame.
• Crackles/explodes at the end.
Equation: \( 2\text{K (s)} + 2\text{H}_2\text{O (l)} \rightarrow 2\text{KOH (aq)} + \text{H}_2\text{ (g)} \)
2. Sodium (\(\text{Na}\)):
• Floats and moves rapidly across the surface.
• Melts into a shiny, silvery ball (due to the heat released).
• Fizzes rapidly, releasing hydrogen gas.
• Solution turns alkaline (phenolphthalein indicator turns pink).
Equation: \( 2\text{Na (s)} + 2\text{H}_2\text{O (l)} \rightarrow 2\text{NaOH (aq)} + \text{H}_2\text{ (g)} \)
3. Calcium (\(\text{Ca}\)):
• Sinks initially, then rises as bubbles of hydrogen attach to it.
• Fizzes vigorously.
• The water turns cloudy/milky because calcium hydroxide (\(\text{Ca(OH)}_2\)) is only slightly soluble and forms a white precipitate/suspension.
Equation: \( \text{Ca (s)} + 2\text{H}_2\text{O (l)} \rightarrow \text{Ca(OH)}_2\text{ (aq/s)} + \text{H}_2\text{ (g)} \)
4. Magnesium to Copper:
• Magnesium: Reacts extremely slowly with cold water (only a few tiny bubbles over several days).
• Aluminium, Zinc, Iron, Copper: Show no reaction with cold water.
Part B: Reactions with Steam
Metals below calcium in the series that do not react well with cold water can react when heated in steam (\(\text{H}_2\text{O (g)}\)). In steam, a metal oxide is produced (not a hydroxide!).
General Word Equation:
\( \text{Metal} + \text{Steam} \rightarrow \text{Metal Oxide} + \text{Hydrogen} \)
• Magnesium (\(\text{Mg}\)): Burns with a bright white flame inside the boiling tube, leaving white magnesium oxide powder:
\( \text{Mg (s)} + \text{H}_2\text{O (g)} \rightarrow \text{MgO (s)} + \text{H}_2\text{ (g)} \)
• Zinc (\(\text{Zn}\)): Glows when heated; produces zinc oxide (yellow when hot, white when cold) and hydrogen gas:
\( \text{Zn (s)} + \text{H}_2\text{O (g)} \rightarrow \text{ZnO (s)} + \text{H}_2\text{ (g)} \)
• Iron (\(\text{Fe}\)): Must be heated strongly; glows red and produces dark tri-iron tetroxide (\(\text{Fe}_3\text{O}_4\)) and hydrogen gas:
\( 3\text{Fe (s)} + 4\text{H}_2\text{O (g)} \rightarrow \text{Fe}_3\text{O}_4\text{ (s)} + 4\text{H}_2\text{ (g)} \)
• Copper (\(\text{Cu}\)): No reaction with steam, even under strong heating.
Experimental Setup and Crucial Safety: "Suck-Back"
To react a metal with steam in the lab, damp mineral wool is placed at the bottom of a horizontal boiling tube to generate steam when warmed, while the metal sample in the middle is heated strongly. Hydrogen gas passes through a delivery tube and is collected over a trough of water.
Safety Warning (Common Exam Question): Always remove the delivery tube from the water trough before turning off the Bunsen burner. If you remove the heat first, the gas inside the tube cools and contracts, creating a vacuum that pulls cold water up the tube ("suck-back"), causing the hot glass tube to crack or shatter!
Section Key Takeaway:
Cold Water \(\rightarrow\) Metal Hydroxide + Hydrogen (\(\text{H}_2\))
Steam \(\rightarrow\) Metal Oxide + Hydrogen (\(\text{H}_2\))
4. Displacement Reactions & Redox
The Displacement Rule
A more reactive metal will displace (take the place of) a less reactive metal from an aqueous solution of its salt or from its solid metal oxide.
Analogy: Think of a stronger player taking the ball away from a weaker player. If the metal on its own is higher up the reactivity series, it takes the compound partner!
1. Displacement in Aqueous Solutions
Let's look at adding silver-grey magnesium ribbon into a blue solution of copper(II) sulfate:
Word Equation:
\( \text{Magnesium} + \text{Copper(II) Sulfate} \rightarrow \text{Magnesium Sulfate} + \text{Copper} \)
Full Chemical Equation:
\( \text{Mg (s)} + \text{CuSO}_4\text{ (aq)} \rightarrow \text{MgSO}_4\text{ (aq)} + \text{Cu (s)} \)
Key Observations to Memorise for Exams:
1. The blue colour of the copper sulfate solution gradually fades and turns colourless.
2. The silver-grey magnesium ribbon dissolves / disappears.
3. A red-brown solid (or pink-brown solid) of copper metal is deposited.
4. The temperature of the mixture rises (it is an exothermic reaction).
Step-by-Step: Writing Net Ionic Equations
Examiners frequently ask you to convert full equations into net ionic equations by removing spectator ions.
Step 1: Write the full equation with state symbols:
\( \text{Mg (s)} + \text{CuSO}_4\text{ (aq)} \rightarrow \text{MgSO}_4\text{ (aq)} + \text{Cu (s)} \)
Step 2: Break all soluble aqueous ionic substances into their separate ions:
\( \text{Mg (s)} + \text{Cu}^{2+}\text{ (aq)} + \text{SO}_4^{2-}\text{ (aq)} \rightarrow \text{Mg}^{2+}\text{ (aq)} + \text{SO}_4^{2-}\text{ (aq)} + \text{Cu (s)} \)
Step 3: Cross out spectator ions (ions that do not change state or charge):
Sulfate (\(\text{SO}_4^{2-}\)) is identical on both sides, so we cross it out.
Step 4: Write the remaining net ionic equation:
\( \text{Mg (s)} + \text{Cu}^{2+}\text{ (aq)} \rightarrow \text{Mg}^{2+}\text{ (aq)} + \text{Cu (s)} \)
Understanding the Redox Aspect
Remember OIL RIG (Oxidation Is Loss of electrons, Reduction Is Gain of electrons):
• Magnesium is oxidised because each atom loses two electrons:
\( \text{Mg} \rightarrow \text{Mg}^{2+} + 2e^- \)
• Copper(II) ions are reduced because each ion gains two electrons:
\( \text{Cu}^{2+} + 2e^- \rightarrow \text{Cu} \)
2. Solid State Displacement: The Thermite Reaction
Displacement also happens between solids when heated strongly. A famous example is the competition between aluminium powder and iron(III) oxide:
\( 2\text{Al (s)} + \text{Fe}_2\text{O}_3\text{ (s)} \rightarrow \text{Al}_2\text{O}_3\text{ (s)} + 2\text{Fe (l/s)} \)
Because aluminium is much higher in the reactivity series than iron, it aggressively removes the oxygen from iron(III) oxide. This reaction releases so much heat that the displaced iron is produced in a molten liquid state, making it ideal for welding railway tracks together on site!
Section Key Takeaway: In displacement reactions, the more reactive metal loses electrons (oxidised) and forces the less reactive metal ion to gain electrons (reduced).
---5. Prescribed Practical C5: Investigating Metal Reactivity
Overview & Method
In Prescribed Practical C5, you investigate and verify the reactivity series by carrying out displacement reactions in a spotting tile or small test tubes.
1. Small pieces of four metals—Magnesium (\(\text{Mg}\)), Zinc (\(\text{Zn}\)), Iron (\(\text{Fe}\)), and Copper (\(\text{Cu}\))—are placed into separate wells or tubes.
2. Solutions of their salts are added: Magnesium nitrate (\(\text{Mg(NO}_3)_2\)), Zinc sulfate (\(\text{ZnSO}_4\)), Iron(II) sulfate (\(\text{FeSO}_4\)), and Copper(II) sulfate (\(\text{CuSO}_4\)).
3. You observe any temperature change (using a thermometer) and visual signs of reaction (colour changes of solutions, deposits of displaced metal).
Summary of Results Table
• Magnesium: Reacts with \(\text{ZnSO}_4\), \(\text{FeSO}_4\), and \(\text{CuSO}_4\) (Displaces all three; largest temperature rise).
• Zinc: Reacts with \(\text{FeSO}_4\) and \(\text{CuSO}_4\), but not \(\text{Mg(NO}_3)_2\).
• Iron: Reacts only with \(\text{CuSO}_4\).
• Copper: Shows no reaction with any of the solutions (displaces none).
Conclusion: By counting the number of reactions, we confirm the order of reactivity: \( \text{Mg} > \text{Zn} > \text{Fe} > \text{Cu} \).
---6. Common Exam Traps & How to Avoid Them
Make sure you don't lose easy marks by keeping these examiner-reported pitfalls in mind:
Trap 1: Writing "Clear" instead of "Colourless"
A copper sulfate solution is blue and clear (transparent). When the reaction finishes, it becomes colourless. "Clear" is not a colour! Always write "blue solution turns colourless".
Trap 2: Misidentifying Copper Metal's Colour
When copper is displaced from solution, it is finely divided and looks red-brown, pink-brown, or brown solid. Never write that copper is "red" or "black".
Trap 3: Mixing up Products of Cold Water vs. Steam
Magnesium with cold water produces magnesium hydroxide: \( \text{Mg(OH)}_2 \).
Magnesium with steam produces magnesium oxide: \( \text{MgO} \).
Trap 4: Leaving Spectator Ions in Ionic Equations
If a question asks specifically for an ionic equation, you must cancel out spectator ions like \( \text{SO}_4^{2-} \) or \( \text{NO}_3^- \). Writing \( \text{Mg} + \text{CuSO}_4 \rightarrow \text{MgSO}_4 + \text{Cu} \) will lose marks if an ionic equation was requested!
Trap 5: Explaining Aluminium's Reactivity
If asked why aluminium does not corrode or react as quickly as expected, always credit the impermeable/protective aluminium oxide (\(\text{Al}_2\text{O}_3\)) layer on its surface.
Quick Revision Checklist
Before moving on, check that you can confidently:
• State the reactivity series from potassium down to copper.
• Explain reactivity in terms of electron loss and ion formation (\(\text{M} \rightarrow \text{M}^{n+} + ne^-\)).
• Recall the flame colours for \(\text{K}\) (lilac), \(\text{Na}\) (yellow/orange), \(\text{Ca}\) (brick-red), and \(\text{Mg}\) (blinding white).
• Write balanced symbol equations for metals reacting with cold water and steam.
• Predict whether a displacement reaction will occur between a given metal and salt solution.
• Construct balanced net ionic equations and explain displacement using oxidation and reduction.