Metals and Non-metals
CBSE Class 10 · Science · Notes, formulas and practice questions
This chapter is about the differences in physical and chemical properties between metals and non-metals, the reactivity series that explains reactions and guides extraction, the nature and properties of ionic compounds, and methods of extracting metals from ores and preventing corrosion.
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Metals and non-metals differ in physical appearance, hardness, conductivity, and chemical reactivity. Metals such as iron and aluminium are lustrous, malleable, ductile, and good conductors of heat and electricity. Non-metals like sulphur and carbon are generally dull, brittle, and poor conductors. These differences arise from their electronic structure: metals tend to lose electrons to form positive ions, while non-metals gain electrons to form negative ions. This fundamental difference underlies all the chemistry in this chapter.
The reactivity series arranges metals in decreasing order of their tendency to lose electrons. Potassium and sodium are at the top, while gold and silver are at the bottom. This series is the key to predicting displacement reactions and choosing an extraction method. A more reactive metal can displace a less reactive metal from its salt solution. In nature, metals that lie high in the series are never found free and must be extracted by electrolysis, whereas lower metals may occur native or require only simple heating. The series thus connects reactivity to practical metallurgy.
Ionic bonding happens when a metal transfers its valence electrons to a non-metal, creating oppositely charged ions. The electrostatic force holding these ions together is called an ionic bond. Ionic compounds form a regular crystalline lattice, which gives them high melting and boiling points. They are brittle solids in the dry state and do not conduct electricity, but they conduct when melted or dissolved because the ions become free to move. Most ionic compounds are soluble in water and insoluble in organic liquids.
Metals occur in the Earth's crust combined with oxygen, sulphur, carbonates, and other substances, as ores. Extraction involves concentrating the ore, converting it to an oxide by roasting or calcination, reducing it to the free metal, and purifying it by electrolytic refining. Corrosion is the undesirable oxidative damage of metals, shown clearly by rusting of iron. Protection is achieved by excluding air and moisture through painting, oiling, galvanising with a more reactive metal, or by making alloys that resist corrosion. Together, these processes describe the full journey from ore to finished metal object.
Key terms
- Reactivity series
- A list of metals arranged in decreasing order of their chemical reactivity, starting from the most reactive (potassium) down to the least reactive (gold). It is used to predict whether a metal can displace another from its salt solution and to decide the method of extraction.
- Ionic bond
- The electrostatic force of attraction between positively charged metal ions and negatively charged non-metal ions, formed by the complete transfer of one or more electrons from the metal to the non-metal. For example, sodium gives an electron to chlorine to form Na⁺ and Cl⁻.
- Roasting
- Heating a sulphide ore strongly in the presence of excess air to convert it into an oxide. For instance, zinc sulphide, ZnS, is roasted to give zinc oxide and sulphur dioxide. Roasting also removes moisture and volatile impurities.
- Calcination
- Heating an ore, usually a carbonate, strongly in a limited supply of air (or in the absence of air) to convert it into the oxide. Calcium carbonate is calcined to produce calcium oxide and carbon dioxide. Calcination also drives off water from hydrated ores.
- Electrolytic refining
- A purification method in which impure metal is made the anode, a thin sheet of pure metal is the cathode, and a water-soluble salt of the metal is the electrolyte. On passing current, pure metal deposits on the cathode while impurities collect as anode mud below the anode.
- Corrosion
- The gradual eating away of a metal through chemical reaction with its surroundings. Rusting of iron is the most common example. Corrosion requires both oxygen and moisture and reduces the strength and life of metal objects.
- Galvanisation
- The process of coating a layer of zinc on iron or steel to protect it from rusting. Zinc acts as a sacrificial metal because it is more reactive than iron, so zinc reacts first in air and moisture, shielding the iron underneath.
- Alloy
- A homogeneous mixture of a metal with one or more other metals or non-metals, made to improve properties such as strength, hardness or resistance to corrosion. Steel, brass and bronze are common alloys. Steel (iron with a little carbon) is much harder than pure iron, while stainless steel (iron with nickel and chromium) is the one that does not rust.
- Anodising
- An electrolytic process used to create a thick and hard oxide layer on the surface of aluminium articles. The aluminium article is made the anode in an electrolyte, and the oxide layer formed protects the aluminium from further corrosion and makes it more durable.
Formula sheet
| What | Formula | Notes |
|---|---|---|
| Metal with oxygen | 2Mg + O₂ → 2MgO | A metal burning in oxygen forms a metal oxide. Metal oxides are basic in nature. |
| Metal with a dilute acid | Zn + 2HCl → ZnCl₂ + H₂ | A metal above hydrogen in the reactivity series displaces hydrogen from a dilute acid, giving a salt and hydrogen gas. |
| Metal with dilute sulphuric acid | Mg + H₂SO₄ → MgSO₄ + H₂ | The same displacement with sulphuric acid instead of hydrochloric acid, producing the sulphate salt. |
| Roasting | 2ZnS + 3O₂ → 2ZnO + 2SO₂ | Heating a sulphide ore in excess air to convert it to the oxide, which is easier to reduce to the metal. |
| Calcination | CaCO₃ → CaO + CO₂ | Heating a carbonate ore in a limited supply of air to convert it to the oxide. |
| Self-reduction in copper extraction | Cu₂S + 2Cu₂O → 6Cu + SO₂ | Copper(I) sulphide reduces copper(I) oxide without any separate reducing agent, which is why copper is extracted this way. |
Practice questions with answers
1. What is the reactivity series? State its importance in the extraction of metals.
The reactivity series is the arrangement of metals in decreasing order of their chemical reactivity, from most reactive potassium to least reactive gold. It helps predict displacement reactions and tells us that highly reactive metals like sodium and calcium must be extracted by electrolysis of their molten ores, while moderately reactive metals like zinc and iron can be reduced using carbon, and less reactive metals like gold may be found free in nature.
2. Why do ionic compounds have high melting and boiling points?
Ionic compounds consist of oppositely charged ions held together by strong electrostatic forces of attraction in a regular lattice. A large amount of heat energy is required to break this lattice apart before the substance can melt or boil. Hence, ionic compounds have high melting and boiling points.
3. Give a reason: Sodium metal is stored in kerosene.
Sodium is a highly reactive metal that reacts violently with oxygen and water vapour present in the air at room temperature. Storing it under kerosene keeps it completely isolated from air and moisture, preventing the vigorous reaction and possible fire hazard.
4. Write the chemical equation when zinc reacts with dilute hydrochloric acid. Name the products.
Zinc, being more reactive than hydrogen, displaces hydrogen from dilute hydrochloric acid. The equation is Zn + 2HCl → ZnCl₂ + H₂. The products are zinc chloride and hydrogen gas, which bubbles out.
5. How is copper extracted from its sulphide ore? Outline the steps with equations.
Copper sulphide ore (Cu₂S) is first roasted in air to convert a part of it to copper oxide, and the remaining sulphide then reacts with this oxide to give copper and sulphur dioxide. The overall reaction is Cu₂S + 2Cu₂O → 6Cu + SO₂. The crude copper is further purified by electrolytic refining.
6. Distinguish between roasting and calcination with examples.
Roasting is heating a sulphide ore in excess air to convert it to oxide; for example, 2ZnS + 3O₂ → 2ZnO + 2SO₂. Calcination is heating a carbonate ore in a limited supply of air to convert it to oxide; for example, CaCO₃ → CaO + CO₂. Roasting produces sulphur dioxide, whereas calcination produces carbon dioxide.
7. What conditions are necessary for the rusting of iron? Explain any one method to prevent rusting.
Rusting of iron requires the presence of both oxygen and water (moisture). Salt and acids speed up the process. One method of prevention is galvanisation, in which a layer of zinc is applied over iron. Zinc is more reactive than iron, so it corrodes preferentially, keeping the iron protected as long as the zinc layer remains intact.
8. What is an alloy? Why are alloys generally preferred over the pure metal? Give one example.
An alloy is a homogeneous mixture of a metal with one or more other metals or non-metals. Alloys are made because they are usually harder and stronger than the pure metal, and some of them also resist corrosion. Steel, for example, is iron alloyed with a small amount of carbon and is far harder than pure iron, which is why it is used in construction; stainless steel, made by alloying iron with nickel and chromium, is the alloy that does not rust.
9. Describe the process of electrolytic refining of copper with a labelled diagram (in words).
Electrolytic refining uses an electrolytic cell with impure copper as the anode, a thin strip of pure copper as the cathode, and an acidified solution of copper sulphate as the electrolyte. When a low-voltage current passes, copper dissolves from the anode into the solution and deposits on the cathode. Impurities such as silver and gold settle as anode mud below the anode, yielding very pure copper at the cathode.
10. State the reactions that occur when a metal like magnesium reacts with oxygen and with dilute sulphuric acid.
Magnesium burns in oxygen to form magnesium oxide: 2Mg + O₂ → 2MgO. With dilute sulphuric acid, magnesium produces magnesium sulphate and hydrogen gas: Mg + H₂SO₄ → MgSO₄ + H₂. In both cases, magnesium loses electrons and becomes a positive ion, showing its metallic nature.
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