Chemical Properties & Reactivity Series - Tenth Grade Science
While physical properties describe the mechanical and thermodynamic states of matter, the chemical properties of metals and non-metals reveal their fundamental tendencies to gain, lose, or share electrons during chemical transformations. Metals are inherently electropositive elements that readily lose valence electrons to form cations, whereas non-metals are electronegative elements that gain or share electrons to achieve stable octet configurations. However, the chemical reactivity of metals varies across orders of magnitude: while potassium ignites explosively upon contact with cold water, gold remains uncorroded over millennia. The reactivity series organizes metals into an empirical hierarchy based on their standard oxidation potentials. In this chapter, you will investigate how metals react with oxygen, water, acids, and salt solutions, explore amphoteric oxides, analyze non-metal chemical dynamics, and master the predictive power of the reactivity series.
Reactions of Metals with Atmospheric Oxygen
Most metals combine with oxygen to form metallic oxides, which are typically basic in chemical character, though notable amphoteric and acidic exceptions exist.
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| OXIDATION DYNAMICS ACROSS METALS |
+-------------------+-----------------------------------+---------------------------+
| Metal | Reaction with Oxygen | Thermal / Ignition Cond. |
+-------------------+-----------------------------------+---------------------------+
| Potassium, Sodium | Ignites spontaneously at room temp| Vigorous, stored in oil |
| Magnesium | Burns with brilliant white flame | Requires initial heating |
| Iron | Does not burn; filings glow bright| High-heat surface oxide |
| Copper | Surface turns black; no ignition | Sustained strong heating |
| Silver, Gold | Completely unreactive with oxygen | Inert at all temperatures |
+-------------------+-----------------------------------+---------------------------+
Basic Oxides versus Amphoteric Oxides
When basic metallic oxides dissolve in water, they yield alkaline solutions containing hydroxide ions:
Na2O(s) + H2O(l) -> 2NaOH(aq)
K2O(s) + H2O(l) -> 2KOH(aq)
Most metallic oxides, such as copper(II) oxide ($\text{CuO}$) and magnesium oxide ($\text{MgO}$), are insoluble in water but dissolve readily in dilute mineral acids to form salts and water, confirming their basic nature:
MgO(s) + 2HCl(aq) -> MgCl2(aq) + H2O(l)
However, certain metallic oxides display dual chemical behavior, reacting with both strong acids and strong alkalis to produce salt and water. These compounds are designated as amphoteric oxides, exemplified by aluminium oxide ($\text{Al}_2\text{O}_3$) and zinc oxide ($\text{ZnO}$):
Reaction with Acid:
Al2O3(s) + 6HCl(aq) -> 2AlCl3(aq) + 3H2O(l) (Behaves as a Base)
Reaction with Base:
Al2O3(s) + 2NaOH(aq) -> 2NaAlO2(aq) + H2O(l) (Behaves as an Acid: Sodium Aluminate)
Similarly, zinc oxide reacts with hydrochloric acid to form zinc chloride, and with sodium hydroxide to form sodium zincate:
ZnO(s) + 2NaOH(aq) -> Na2ZnO2(aq) + H2O(l) (Sodium Zincate)
Reactions of Metals with Water and Steam
The reaction of metals with water produces metallic hydroxides or oxides along with hydrogen gas, with reaction rates dictated by the position of the metal in the reactivity series.
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| METAL-WATER INTERACTION MATRIX |
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| Metal | State of Water Required | Primary Products Formed |
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| Potassium, Sodium | Violent in cold water | Hydroxide + H2(g) (Ignites) |
| Calcium | Rapid in cold water | Ca(OH)2 + H2(g) (Floats on bubbles|
| Magnesium | Slow in cold; fast in hot | Mg(OH)2 + H2(g) (Floats on bubbles|
| Aluminium, Fe, Zn | Inactive with hot water; | Oxide + H2(g) |
| | reacts with steam | |
| Copper, Ag, Au | Completely unreactive | No reaction |
+-------------------+---------------------------+-----------------------------------+
Alkali and Alkaline Earth Metals with Cold Water
Potassium and sodium react violently and exothermically with cold water:
2K(s) + 2H2O(l) -> 2KOH(aq) + H2(g) + Heat
2Na(s) + 2H2O(l) -> 2NaOH(aq) + H2(g) + Heat
The enthalpy of reaction is so great that the liberated hydrogen gas ignites immediately in air, burning with a characteristic lilac flame for potassium and a golden-yellow flame for sodium. Because sodium and potassium react spontaneously with atmospheric moisture and oxygen, they are stored submerged beneath dry kerosene oil.
Calcium reacts less violently with cold water, forming cloudy slaked lime and hydrogen gas:
Ca(s) + 2H2O(l) -> Ca(OH)2(aq) + H2(g)
The heat evolved is insufficient to ignite the hydrogen. The calcium metal pieces sink initially, then float to the liquid surface as adhering hydrogen gas bubbles provide buoyant lift.
Reaction with Steam
Metals in the middle of the reactivity series—such as aluminium, zinc, and iron—do not react with cold or hot liquid water. They react only with superheated steam, producing metal oxides rather than hydroxides:
2Al(s) + 3H2O(g) -> Al2O3(s) + 3H2(g)
Zn(s) + H2O(g) -> ZnO(s) + H2(g)
3Fe(s) + 4H2O(g) -> Fe3O4(s) + 4H2(g) (Magnetic Iron Oxide)
Reactions with Acids and Aqua Regia
Metals situated above hydrogen in the reactivity series displace hydrogen from dilute hydrochloric and sulfuric acids.
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| REACTION OF METALS WITH MINERAL ACIDS |
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| Mg(s) + 2HCl(aq) -----> MgCl2(aq) + H2(g) (Most vigorous effervescence) |
| Al(s) + 6HCl(aq) -----> 2AlCl3(aq) + 3H2(g) (Induction delay due to oxide) |
| Zn(s) + 2HCl(aq) -----> ZnCl2(aq) + H2(g) (Moderate effervescence) |
| Fe(s) + 2HCl(aq) -----> FeCl2(aq) + H2(g) (Slow bubble evolution) |
| Cu(s) + HCl(aq) -----> No Reaction (Below hydrogen in series) |
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The Exceptional Behavior of Nitric Acid
Dilute nitric acid ($\text{HNO}_3$) typically does not evolve hydrogen gas when reacted with metals. Nitric acid is a powerful oxidizing agent; it oxidizes the evolving hydrogen gas to water while itself undergoing reduction to nitrogen dioxide ($\text{NO}_2$), nitric oxide ($\text{NO}$), or nitrous oxide ($\text{N}_2\text{O}$). However, very dilute ($1\%$) nitric acid reacts with magnesium ($\text{Mg}$) and manganese ($\text{Mn}$) to liberate hydrogen gas, as their high oxidation potentials overcome the acid's oxidizing capacity:
Mg(s) + 2HNO3(aq) [Very Dilute] -> Mg(NO3)2(aq) + H2(g)
The Dissolving Power of Aqua Regia
Noble metals such as gold ($\text{Au}$) and platinum ($\text{Pt}$) are unaffected by concentrated hydrochloric acid, concentrated nitric acid, or any single mineral acid. However, they dissolve in aqua regia (royal water), a freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in a $3:1$ volume ratio:
HNO3(conc) + 3HCl(conc) -> NOCl(g) + 2H2O(l) + 2[Cl] (Nascent Chlorine)
Aqua regia generates nascent chlorine atoms and nitrosyl chloride ($\text{NOCl}$), which attack noble metals to form soluble coordination complexes:
Au(s) + 3[Cl] + HCl(aq) -> HAuCl4(aq) (Chloroauric Acid)
The Electrochemical Reactivity Series
The reactivity series is an empirical ranking of metals arranged in decreasing order of their chemical reactivity and electropositive character.
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| THE REACTIVITY SERIES |
+-----------------------------------+-----------------------------------------------+
| Potassium (K) | Most Reactive Metal (Displaces all below) |
| Sodium (Na) | Reacts violently with cold water |
| Calcium (Ca) | Extracted by molten electrolysis |
| Magnesium (Mg) | |
| Aluminium (Al) |-----------------------------------------------|
| Zinc (Zn) | Moderately Reactive Metals |
| Iron (Fe) | React with steam and mineral acids |
| Lead (Pb) | Extracted by carbon reduction |
| [HYDROGEN] (H) | Non-metal reference datum (E0 = 0.00 V) |
| Copper (Cu) |-----------------------------------------------|
| Mercury (Hg) | Least Reactive Metals |
| Silver (Ag) | Found native or extracted by thermal roasting |
| Gold (Au) | Unreactive with water, acids, and oxygen |
+-----------------------------------+-----------------------------------------------+
Displacement Criteria in Aqueous Solutions
Any metal situated higher in the reactivity series will displace a metal situated below it from an aqueous solution of its salt. For example, zinc displaces copper from copper sulfate solution:
Zn(s) + CuSO4(aq) -> ZnSO4(aq) + Cu(s)
Because zinc has a higher tendency to lose electrons, it oxidizes to $\text{Zn}^{2+}$, while copper ions accept electrons to precipitate as metallic copper. Conversely, if copper is added to zinc sulfate solution, no reaction occurs.
Chemical Properties of Non-Metals
Non-metals exhibit electronegative behavior, acting as electron acceptors and oxidizing agents.
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| CHEMICAL PATTERNS OF NON-METALS |
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| 1. Formation of Acidic Oxides: C + O2 -> CO2; CO2 + H2O -> H2CO3 |
| S + O2 -> SO2; SO2 + H2O -> H2SO3 |
| 2. Formation of Neutral Oxides: 2C + O2 -> 2CO; 2H2 + O2 -> 2H2O; N2O |
| 3. Inactivity with Dilute Acid: C, S, P do not displace H2 from mineral acids. |
| 4. Covalent Chlorides: P4 + 6Cl2 -> 4PCl3 (Liquid, hydrolyzed by water)|
| 5. Covalent Hydrides: N2 + 3H2 -> 2NH3; H2 + S -> H2S |
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Acidic and Neutral Oxides of Non-Metals
Non-metals react with oxygen to produce covalent oxides that are acidic or neutral. Acidic oxides dissolve in water to yield mineral acids:
SO3(g) + H2O(l) -> H2SO4(aq) (Sulfuric Acid)
Neutral non-metallic oxides, including carbon monoxide ($\text{CO}$), water ($\text{H}_2\text{O}$), and nitrous oxide ($\text{N}_2\text{O}$), do not exhibit acidic or basic properties and do not alter litmus paper.
Non-metals do not displace hydrogen gas from dilute acids because non-metals are electron acceptors and cannot reduce hydrogen ions ($\text{H}^+$) into elemental hydrogen gas.
Chapter Exercises
Exercise 1: An element A burns in air with a golden-yellow flame to form an oxide B. When B is dissolved in water, it forms an alkaline solution C that turns red litmus paper blue and produces hydrogen gas when boiled with zinc metal. Identify A, B, and C, and write balanced chemical equations for all described reactions.
Exercise 2: Write balanced chemical equations demonstrating that zinc oxide is an amphoteric oxide by showing its reactions with dilute sulfuric acid and with concentrated aqueous potassium hydroxide.
Exercise 3: Explain why copper does not evolve hydrogen gas when placed in dilute hydrochloric acid, but dissolves in concentrated nitric acid with the evolution of reddish-brown fumes.
Solutions and Explanations
Solution 1: Element A is metallic sodium ($\text{Na}$), which burns in air with a characteristic golden-yellow flame to form basic sodium oxide ($\text{Na}_2\text{O}$), compound B: $4\text{Na(s)} + \text{O}_2\text{(g)} \rightarrow 2\text{Na}_2\text{O(s)}$. When sodium oxide dissolves in water, it produces sodium hydroxide ($\text{NaOH}$), compound C: $\text{Na}_2\text{O(s)} + \text{H}_2\text{O(l)} \rightarrow 2\text{NaOH(aq)}$. Sodium hydroxide is an alkali that turns red litmus paper blue. When boiled with zinc metal, sodium hydroxide reacts to form soluble sodium zincate and hydrogen gas: $\text{Zn(s)} + 2\text{NaOH(aq)} \rightarrow \text{Na}_2\text{ZnO}_2\text{(aq)} + \text{H}_2\text{(g)}$.
Solution 2: Zinc oxide ($\text{ZnO}$) behaves as an amphoteric oxide by reacting with both acids and alkalis. Reacting with dilute sulfuric acid, zinc oxide acts as a basic oxide, neutralizing the acid to form zinc sulfate and water: $\text{ZnO(s)} + \text{H}_2\text{SO}_4\text{(aq)} \rightarrow \text{ZnSO}_4\text{(aq)} + \text{H}_2\text{O(l)}$. Reacting with concentrated aqueous potassium hydroxide, zinc oxide acts as an acidic oxide, forming potassium zincate and water: $\text{ZnO(s)} + 2\text{KOH(aq)} \rightarrow \text{K}_2\text{ZnO}_2\text{(aq)} + \text{H}_2\text{O(l)}$.
Solution 3: Copper is situated below hydrogen in the electrochemical reactivity series ($E^\circ_{\text{Cu}^{2+}/\text{Cu}} = +0.34\text{ V}$ versus $E^\circ_{\text{H}^+/\text{H}_2} = 0.00\text{ V}$), meaning copper has a lower oxidation potential than hydrogen and cannot reduce hydronium ions in dilute hydrochloric acid to liberate hydrogen gas. However, concentrated nitric acid is a potent oxidizing agent. Rather than functioning as a typical acid-base displacement, concentrated nitric acid oxidizes metallic copper to copper(II) ions while the nitrogen atom in the nitrate ion is reduced from $+5$ to $+4$, evolving brown fumes of nitrogen dioxide gas: $\text{Cu(s)} + 4\text{HNO}_3\text{(conc)} \rightarrow \text{Cu(NO}_3)_2\text{(aq)} + 2\text{NO}_2\text{(g)} + 2\text{H}_2\text{O(l)}$.