👋 Welcome to the World of Reactivity!
Hello future chemist! This chapter is incredibly important because it explains why some metals are stored under oil while others sit happily as jewellery on your finger.
We are going to learn about the Reactivity Series—a simple ranking system that predicts how different metals behave when they meet water, acids, or other metal compounds. We will also explore displacement, redox reactions, and the chemistry of rusting!
🔬 Section 1: What Exactly is Reactivity?
In chemistry, the reactivity of a metal refers to how easily that metal reacts, which corresponds to how easily it loses its outer electrons to form positive ions (cations).
- A highly reactive metal (like Potassium or Sodium) loses its outer electrons very easily to form positive ions.
- A low reactivity metal (like Gold or Silver) holds onto its electrons tightly and is unreactive.
The Key Takeaway: The higher a metal is in the reactivity series, the more easily it forms positive ions and the more vigorous its reactions are.
Quick Review Box: Reactivity vs. Electrons
High Reactivity = Readily loses electrons to form positive ions
Low Reactivity = Reluctant to lose electrons (Relatively unreactive/stable)
💡 Section 2: The Order of the Reactivity Series
The reactivity series arranges metals in order of decreasing reactivity. Non-metals like Carbon and Hydrogen are often included for comparison (useful for metal extraction and acid reactions).
The Official Edexcel IGCSE Reactivity Series
- Potassium ( ext{K}) - Most reactive
- Sodium ( ext{Na})
- Lithium ( ext{Li})
- Calcium ( ext{Ca})
- Magnesium ( ext{Mg})
- Aluminium ( ext{Al})
- (Carbon, ext{C})
- Zinc ( ext{Zn})
- Iron ( ext{Fe})
- (Hydrogen, ext{H})
- Copper ( ext{Cu})
- Silver ( ext{Ag})
- Gold ( ext{Au}) - Least reactive
🚀 Memory Aid: Mnemonic!
Use this phrase to remember the order of the metals:
Please Send Lions, Cats, Monkeys, And Zebras Into Cool Shiny Gardens.
(Potassium, Sodium, Lithium, Calcium, Magnesium, Aluminium, Zinc, Iron, Copper, Silver, Gold)
🌊 Section 3: Reactivity with Water and Acids
We can deduce the reactivity order by observing how metals react with water and dilute hydrochloric or sulfuric acid.
3.1 Reactions with Water
-
Potassium, Sodium, Lithium, and Calcium: React vigorously with cold water to form a metal hydroxide and hydrogen gas.
\(2\text{Na} (\text{s}) + 2\text{H}_2\text{O} (\text{l}) \rightarrow 2\text{NaOH} (\text{aq}) + \text{H}_2 (\text{g})\) -
Magnesium, Zinc, and Iron: React very slowly with cold water, but react with steam to form a metal oxide and hydrogen gas.
\(\text{Mg} (\text{s}) + \text{H}_2\text{O} (\text{g}) \rightarrow \text{MgO} (\text{s}) + \text{H}_2 (\text{g})\) - Copper, Silver, and Gold: Do not react with cold water or steam.
- Apparent Unreactivity of Aluminium: Aluminium does not react immediately with water or dilute acid because it has a tough, protective aluminium oxide (\(\text{Al}_2\text{O}_3\)) layer on its surface that prevents reaction until removed.
3.2 Reactions with Dilute Acids (\(\text{HCl}\) and \(\text{H}_2\text{SO}_4\))
General Rule: \(\text{Metal} + \text{Acid} \rightarrow \text{Salt} + \text{Hydrogen Gas}\)
- Potassium, Sodium, Lithium: React explosively and violently (too dangerous for school laboratory testing).
- Calcium, Magnesium, Aluminium, Zinc, Iron: React steadily with dilute hydrochloric or sulfuric acid, producing bubbles of \(\text{H}_2\) gas and forming a solution of the metal chloride or sulfate.
\(\text{Zn} (\text{s}) + 2\text{HCl} (\text{aq}) \rightarrow \text{ZnCl}_2 (\text{aq}) + \text{H}_2 (\text{g})\)
\(\text{Mg} (\text{s}) + \text{H}_2\text{SO}_4 (\text{aq}) \rightarrow \text{MgSO}_4 (\text{aq}) + \text{H}_2 (\text{g})\) - Copper, Silver, Gold: Do not react with dilute acids as they are below hydrogen in the reactivity series.
🔄 Section 4: Displacement Reactions
A displacement reaction occurs when a more reactive metal displaces (takes the place of) a less reactive metal from its oxide or from an aqueous solution of its salt.
4.1 Displacement in Aqueous Salt Solutions
When iron is placed into copper(II) sulfate solution, iron displaces copper because iron is more reactive than copper:
\(\text{Fe} (\text{s}) + \text{CuSO}_4 (\text{aq}) \rightarrow \text{FeSO}_4 (\text{aq}) + \text{Cu} (\text{s})\)
Observations: The blue solution turns pale green (as \(\text{Cu}^{2+}\) ions are replaced by \(\text{Fe}^{2+}\) ions), and a pink/red-brown deposit of copper coats the iron.
4.2 Displacement with Solid Metal Oxides
Heating a reactive metal with the oxide of a less reactive metal causes displacement:
\(\text{Mg} (\text{s}) + \text{CuO} (\text{s}) \rightarrow \text{MgO} (\text{s}) + \text{Cu} (\text{s})\)
⚡ Section 5: Oxidation, Reduction, and Redox
Displacement reactions are examples of redox reactions, where oxidation and reduction occur simultaneously.
Definitions in Terms of Oxygen:
- Oxidation: Gain of oxygen.
- Reduction: Loss of oxygen.
- Oxidising agent: A substance that gives oxygen to another substance (and is itself reduced).
- Reducing agent: A substance that takes oxygen away from another substance (and is itself oxidised).
Example: \(\text{CuO} + \text{Mg} \rightarrow \text{Cu} + \text{MgO}\)
\(\text{CuO}\) is reduced (loses oxygen) and acts as the oxidising agent.
\(\text{Mg}\) is oxidised (gains oxygen) and acts as the reducing agent.
Definitions in Terms of Electrons (OIL RIG):
- Oxidation Is Loss of electrons (\(\text{OIL}\)).
- Reduction Is Gain of electrons (\(\text{RIG}\)).
- An oxidising agent gains electrons; a reducing agent loses electrons.
Ionic analysis for displacement: \(\text{Fe} (\text{s}) + \text{Cu}^{2+} (\text{aq}) \rightarrow \text{Fe}^{2+} (\text{aq}) + \text{Cu} (\text{s})\)
\(\text{Fe} \rightarrow \text{Fe}^{2+} + 2\text{e}^-\) (Oxidation — \(\text{Fe}\) is the reducing agent)
\(\text{Cu}^{2+} + 2\text{e}^- \rightarrow \text{Cu}\) (Reduction — \(\text{Cu}^{2+}\) is the oxidising agent)
🧱 Section 6: Rusting of Iron and Prevention
Rusting is the corrosion of iron (and steel) forming hydrated iron(III) oxide.
6.1 Conditions Required for Rusting
Both oxygen and water must be present for iron to rust.
- Iron in dry air (using anhydrous calcium chloride to absorb water) \(\rightarrow\) does not rust.
- Iron in boiled water covered with an oil layer (no dissolved oxygen) \(\rightarrow\) does not rust.
- Iron in both air and water \(\rightarrow\) rusts.
6.2 Methods of Rust Prevention
- Barrier Methods: Coating iron with paint, grease/oil, or plastic. This keeps out water and oxygen. If the barrier is scratched, the iron underneath will rust.
- Galvanising: Coating iron with a layer of zinc. Even if the zinc layer is scratched, the zinc protects the iron because zinc is more reactive than iron and reacts preferentially.
- Sacrificial Protection: Attaching blocks of a more reactive metal (such as zinc or magnesium) to iron structures (e.g., ships' hulls, underground pipes). The more reactive metal oxidises/corrodes sacrificially instead of the iron.
✅ Chapter Summary: Quick Review
- The Series: \(\text{K} > \text{Na} > \text{Li} > \text{Ca} > \text{Mg} > \text{Al} > [\text{C}] > \text{Zn} > \text{Fe} > [\text{H}] > \text{Cu} > \text{Ag} > \text{Au}\).
- Water & Acids: High-reactivity metals react with cold water; medium-reactivity metals react with steam and dilute acids; low-reactivity metals below hydrogen do not react.
- Displacement: A more reactive metal displaces a less reactive metal from its salt solution or oxide.
- Redox: Oxidation is gain of oxygen / loss of electrons; reduction is loss of oxygen / gain of electrons.
- Rusting: Requires oxygen and water; prevented by barrier methods, galvanising, and sacrificial protection.