Reactivity series Revision Notes

The reactivity series is a list of metals arranged in order of decreasing reactivity. Reactivity here refers to how readily a metal loses electrons to form positive ions, which determines how easily it reacts with substances such as water, acids, and other metal ions.

This series helps chemists predict the products and feasibility of reactions involving metals, such as displacement reactions and reactions with acids or water. Carbon is included in the series as a reference point because it is used in metal extraction processes.

• Reactivity is related to the ease of electron loss (oxidation) by metals.
• The series ranks metals from most reactive to least reactive.
• Carbon is included to compare metal reactivity with a non-metal reducing agent.

The reactivity series relevant to this objective includes the following elements in order of decreasing reactivity: potassium, sodium, calcium, magnesium, aluminium, carbon, zinc, iron, hydrogen, copper, silver, and gold.

Potassium is the most reactive metal in this list, while gold is the least reactive. Carbon is placed between aluminium and zinc to indicate its role in reducing metal oxides during extraction.

• Potassium (K)
• Sodium (Na)
• Calcium (Ca)
• Magnesium (Mg)
• Aluminium (Al)
• Carbon (C) (non-metal)
• Zinc (Zn)
• Iron (Fe)   Hydrogen (H)   Copper (Cu)   Silver (Ag)   Gold (Au)

Metals higher in the reactivity series lose electrons more easily to form positive ions. This means they react more vigorously with water and acids, producing hydrogen gas and metal compounds.

For example, potassium reacts explosively with cold water, while gold does not react under normal conditions.

The position of hydrogen in the series is important because metals above hydrogen can displace hydrogen from dilute acids, producing hydrogen gas, whereas metals below hydrogen cannot.

• Reactivity depends on ease of electron loss (oxidation).
• Metals above hydrogen react with dilute acids to produce hydrogen gas.
• Metals below hydrogen do not react with dilute acids to release hydrogen.

Potassium, sodium, and calcium react vigorously with cold water to form hydroxides and hydrogen gas. For example:

2K(s) + 2H2O(l) → 2KOH(aq) + H2(g) (potassium)  [Highly vigorous, produces lilac flame]

Magnesium reacts slowly with cold water but more rapidly with steam to form magnesium oxide and hydrogen:  Mg(s) + H2O(g) → MgO(s) + H2(g)  [White solid formed]  [Requires steam]  [Magnesium is less reactive than calcium]  [Not in this objective but relevant for context]  Metals like zinc and iron react with dilute acids to produce salts and hydrogen gas, e.g.:  Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)  [Zinc is above hydrogen]  Copper, silver, and gold do not react with dilute acids under normal conditions because they are less reactive than hydrogen.

• Potassium reacts explosively with cold water.
• Magnesium reacts with steam, not cold water, to form oxide and hydrogen.
• Zinc and iron react with dilute acids to produce hydrogen.
• Copper, silver, and gold do not react with dilute acids.

A metal higher in the reactivity series can displace a metal lower in the series from its compound in solution. For example, zinc can displace copper from copper sulfate solution:

Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)  [Zinc is more reactive than copper]  [Displacement reaction]  [Used to extract or purify metals]  Conversely, copper cannot displace zinc from zinc sulfate solution because it is less reactive.

This principle is used to predict whether a reaction will occur between a metal and a solution of another metal's ions.

• More reactive metals displace less reactive metals from solutions.
• Displacement reactions confirm the order of the reactivity series.
• Example: zinc displaces copper from copper sulfate solution.

The reactivity series is a list of metals arranged in order of decreasing reactivity. It predicts how metals react with water and acids. Metals higher in the series react more vigorously, while those lower down are less reactive or unreactive under similar conditions.

Reactivity depends on how easily a metal loses electrons to form positive ions (oxidation). Metals that lose electrons more readily react faster and more vigorously with water and acids.

• Potassium, sodium, and calcium are very reactive metals at the top of the series.
• Magnesium, zinc, and iron are moderately reactive metals.
• Copper, silver, and gold are low-reactivity metals near the bottom of the series.

Potassium, sodium, and calcium react readily with cold water to form metal hydroxides and hydrogen gas. These reactions are vigorous and often exothermic.

Observations include fizzing or bubbling due to hydrogen gas, metal dissolving or moving on the water surface, and colour changes in the solution due to hydroxide formation.

The general equation for these reactions is: M (s) + H₂O (l) → M⁺ (aq) + OH⁻ (aq) + H₂ (g), where M is the metal.

• Potassium reacts very vigorously, often igniting the hydrogen gas with a lilac flame.
• Sodium reacts vigorously, melting into a ball and fizzing.
• Calcium reacts less vigorously but still produces bubbles of hydrogen.
• All form alkaline solutions due to metal hydroxide formation.

Magnesium does not react significantly with cold water but reacts with steam (hot water vapour) to form magnesium oxide and hydrogen gas.

This reaction requires heating to produce steam and is less vigorous than reactions of potassium or sodium with cold water.

The balanced equation is: Mg (s) + H₂O (g) → MgO (s) + H₂ (g).

• White magnesium oxide forms as a solid product.
• Hydrogen gas is released, observed as bubbles or a squeaky pop when tested.
• The reaction demonstrates magnesium's moderate reactivity, requiring steam rather than cold water.

Magnesium, zinc, and iron react with dilute hydrochloric acid to produce hydrogen gas and their respective metal chlorides in aqueous solution.

These reactions are typical acid-metal reactions where the metal is oxidised and hydrogen ions are reduced.

General equation: Metal (s) + 2HCl (aq) → Metal chloride (aq) + H₂ (g).

• Magnesium reacts rapidly, producing bubbles of hydrogen and dissolving quickly.
• Zinc reacts moderately, with visible fizzing and gradual dissolution.
• Iron reacts slowly, with less vigorous bubbling.
• The solutions formed are colourless metal chloride solutions (e.g., MgCl₂, ZnCl₂, FeCl₂ or FeCl₃ depending on conditions).

Copper, silver, and gold do not react with dilute hydrochloric acid because they are less reactive than hydrogen and cannot displace hydrogen ions from the acid.

No visible reaction occurs; no hydrogen gas is produced and the metals remain unchanged.

This lack of reaction reflects their low position in the reactivity series and their reluctance to lose electrons to form positive ions under these conditions.

• No bubbles or fizzing observed.
• Metal surface remains unchanged.
• These metals require stronger oxidising agents or different conditions to react.

Metal reactivity refers to the tendency of a metal to lose electrons and form positive ions (cations). This electron loss is an oxidation process, and metals that lose electrons more readily are considered more reactive.

The reactivity of a metal influences how vigorously it reacts with substances such as water, acids, and other metal salts. More reactive metals react faster and more completely, often producing hydrogen gas and metal compounds.

The ability of a metal to displace another metal from its compound in solution is a key indicator of relative reactivity. A more reactive metal will displace a less reactive metal from its salt solution, demonstrating its greater tendency to form positive ions.

• Reactivity depends on ease of electron loss (oxidation).
• More reactive metals react faster with water and acids.
• Displacement reactions reveal relative reactivity between metals.

Two common experimental approaches to infer metal reactivity are:

1. Reacting metals with dilute acids (e.g., hydrochloric acid) and observing the rate of hydrogen gas evolution.

2. Performing displacement reactions by adding a metal to a solution of another metal salt and observing whether a reaction occurs, indicated by colour changes or solid deposition.

• Use dilute hydrochloric acid to test metal reactivity safely.
• Observe hydrogen gas bubbles and reaction speed as indicators.
• Displacement reactions involve a metal and a metal salt solution.
• A reaction occurs if the metal is more reactive than the metal ion in solution.

When a metal reacts with dilute hydrochloric acid, the general reaction is:

Metal + Acid → Metal salt + Hydrogen gas (bubbles)

For example, magnesium reacting with hydrochloric acid: Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)↑ (bubbles) (balanced) Observations include effervescence (bubbling) and gradual disappearance of the metal if it is reactive enough. Less reactive metals may show no visible reaction under these conditions. In displacement reactions, for example, zinc metal added to copper(II) sulfate solution: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s) Observations include the blue colour of copper sulfate fading and reddish-brown copper metal forming. No reaction occurs if the metal is less reactive than the metal ion in solution.

• Hydrogen gas evolution indicates metal reacting with acid.
• Bubbles and reaction speed correlate with reactivity.
• Colour changes and solid metal deposition indicate displacement.
• Write balanced symbol equations for all reactions.

Apparatus: small pieces of metals (e.g., magnesium, zinc, iron, copper), dilute hydrochloric acid, test tubes, test tube rack, measuring cylinder, stopwatch, safety goggles, gloves.

Method: 1. Add a fixed volume (e.g., 10 cm³) of dilute hydrochloric acid to separate test tubes. 2. Add equal-sized pieces of different metals to each test tube simultaneously. 3. Observe and time the rate of hydrogen gas evolution (bubbling). 4. Record observations such as bubble rate, metal disappearance, and solution colour changes. 5. Repeat the experiment to ensure reliability and calculate average reaction rates.

Variables: Independent variable: type of metal. Dependent variable: rate of hydrogen gas evolution. Controlled variables: acid concentration, volume, temperature, metal surface area. Safety: Wear goggles and gloves; handle acids carefully; perform in a well-ventilated area; dispose of chemicals properly.

• Use equal metal sizes and acid volumes for fair comparison.
• Measure reaction rate by timing bubble formation.
• Repeat trials to improve accuracy.
• Control variables to ensure valid results.
• Wear appropriate safety equipment.

From the recorded observations, metals can be ranked by their reactivity based on:

Speed and vigour of reaction with acid (faster bubbling = more reactive).

Ability to displace other metals from their salt solutions (displacement occurs only if the metal is more reactive). For example, if magnesium reacts vigorously with acid and displaces zinc from zinc sulfate solution, but zinc does not displace magnesium, magnesium is more reactive than zinc. This comparative approach allows construction of a reactivity series from the experimental data, listing metals from most reactive to least reactive based on their chemical behaviour.

• Rank metals by reaction speed with acid.
• Confirm order by displacement reactions.
• Use observations to build a reactivity series.
• Predict reaction outcomes using inferred order.

Metals react chemically by losing electrons to form positive ions (cations). This process is called oxidation. The ease with which a metal atom loses electrons determines its reactivity.

More reactive metals lose electrons more readily, forming positive ions more easily. For example, magnesium atoms lose two electrons to form Mg²⁺ ions:

Mg (s) → Mg²⁺ (aq) + 2e⁻ (oxidation) (in aqueous solution or ionic compounds)

• Reactivity is linked to electron loss (oxidation).
• Formation of positive ions is a key indicator of metal reactivity.

A displacement reaction occurs when a more reactive metal displaces a less reactive metal from its aqueous ionic compound. This reaction involves electron transfer: the more reactive metal is oxidised (loses electrons), and the less reactive metal ion is reduced (gains electrons) to form the elemental metal.

For example, magnesium metal can displace copper from copper sulfate solution:

Mg (s) + CuSO₄ (aq) → MgSO₄ (aq) + Cu (s) (balanced equation)

• Displacement reactions demonstrate relative reactivity.
• More reactive metals displace less reactive metals from solutions of their ions.
• Electron transfer underlies displacement reactions.

Among the metals magnesium, zinc, iron, copper, and silver, the order of reactivity from most to least reactive is:

Magnesium > Zinc > Iron > Copper > Silver. This order is based on their ability to form positive ions and displace each other from solutions.

For example: magnesium displaces zinc, iron, copper, and silver ions; zinc displaces iron, copper, and silver ions; iron displaces copper and silver ions; copper displaces silver ions; silver does not displace any of these metals.

• Magnesium is the most reactive among these metals.
• Silver is the least reactive and does not displace others.
• This order predicts displacement reaction outcomes.

Displacement reactions can be represented by ionic equations showing electron transfer. Examples include:

1. Magnesium displacing copper ions: Mg (s) → Mg²⁺ (aq) + 2e⁻ (oxidation)

Cu²⁺ (aq) + 2e⁻ → Cu (s) (reduction) Overall: Mg (s) + Cu²⁺ (aq) → Mg²⁺ (aq) + Cu (s) (balanced) 2. Zinc displacing iron ions: Zn (s) → Zn²⁺ (aq) + 2e⁻ Fe²⁺ (aq) + 2e⁻ → Fe (s) Overall: Zn (s) + Fe²⁺ (aq) → Zn²⁺ (aq) + Fe (s) 3. Iron displacing copper ions: Fe (s) → Fe²⁺ (aq) + 2e⁻ Cu²⁺ (aq) + 2e⁻ → Cu (s) Overall: Fe (s) + Cu²⁺ (aq) → Fe²⁺ (aq) + Cu (s)

• Ionic equations show oxidation and reduction half-reactions.
• Electrons lost by the metal are gained by the metal ion.
• Equations must be balanced for mass and charge.

During displacement reactions, the solution’s colour changes as metal ions are replaced and elemental metal forms as a solid deposit.

For example, when magnesium is added to copper sulfate solution:

The blue colour of Cu²⁺ ions fades as copper metal forms as a reddish-brown solid on the magnesium surface. The solution becomes colourless as Mg²⁺ ions form, which are colourless in solution. Practical steps include: using clean metal strips, adding excess metal to ensure complete reaction, stirring, and observing colour changes carefully.

• Colour changes indicate displacement reactions.
• Solid metal deposits form on the more reactive metal.
• Careful observation confirms relative reactivity.

Aluminium is a metal that is chemically reactive. It readily loses electrons to form aluminium ions (Al³⁺) in reactions with oxygen and water under suitable conditions. For example, aluminium reacts with oxygen to form aluminium oxide:

2Al (s) + 3O₂ (g) → 2Al₂O₃ (s) (balanced equation)

This reaction is highly exothermic and occurs rapidly when aluminium is finely divided or heated. However, bulk aluminium metal does not appear to react as readily as expected in everyday conditions.

• Aluminium metal tends to lose three electrons to form Al³⁺ ions.
• It reacts with oxygen to form aluminium oxide (Al₂O₃).

When aluminium is exposed to air, it reacts with oxygen to form a very thin layer of aluminium oxide (Al₂O₃) on its surface. This layer is typically only a few nanometres thick but is extremely dense and adherent. The formation of this oxide layer is a rapid process that occurs at room temperature and normal atmospheric pressure.

The aluminium oxide layer is chemically stable and insoluble in water and most acids and alkalis at room temperature. Because it forms a continuous coating, it acts as a barrier preventing oxygen, water vapour, and other corrosive agents from reaching the underlying aluminium metal.

• Aluminium oxide layer forms immediately on exposure to air.
• The oxide layer is thin, dense, and adherent.
• It is chemically stable and insoluble in water and dilute acids at room temperature.

The aluminium oxide layer protects the metal from further oxidation and corrosion by physically blocking contact between the metal and environmental reactants. This prevents the aluminium atoms beneath the oxide from reacting with oxygen or moisture.

If the oxide layer is damaged mechanically, it reforms quickly in air, restoring the protective barrier. This self-healing property contributes to aluminium's resistance to corrosion and its apparent unreactivity in everyday conditions.

This protective effect explains why aluminium does not react visibly with water or air despite being a reactive metal.

• Oxide layer acts as a barrier to oxygen and water.
• Prevents further oxidation and corrosion of aluminium.
• Self-heals if scratched or damaged, maintaining protection.

Although aluminium appears unreactive, this is due to the protective oxide layer rather than a lack of chemical reactivity. When the oxide layer is removed or disrupted under certain conditions, aluminium reacts readily.

For example, powdered aluminium or aluminium foil treated with acid to remove the oxide layer reacts vigorously with water or acids, producing hydrogen gas:

2Al (s) + 6HCl (aq) → 2AlCl₃ (aq) + 3H₂ (g) (balanced equation)

• Aluminium is chemically reactive but protected by oxide layer.
• Removing the oxide layer exposes reactive aluminium metal.
• Aluminium reacts with acids and water once oxide is removed.

Aluminium metal reacts with oxygen to form aluminium oxide (Al₂O₃).

The oxide layer is thin, stable, and adheres strongly to the metal surface.

This layer prevents further oxidation and corrosion by blocking contact with air and moisture, making aluminium appear unreactive in normal conditions.

• Aluminium oxide layer forms rapidly on exposure to air.
• The layer is stable, insoluble, and self-healing.
• This explains aluminium's apparent unreactivity despite its chemical activity.