Metals And Non-Metals

Difference between Metals and Non-Metals

Graphite can conduct electricity, yet it is a non-metal. Aluminium is high in the reactivity series, yet an aluminium object can resist further corrosion. Solid sodium chloride contains charged ions, yet it does not conduct like a copper wire. These apparent contradictions become useful once we ask what the particles and surfaces are doing.

This note follows Class 10 Science, NCERT Chapter 3, Metals and Non-metals, Reprint 2026–27, and the corresponding CBSE curriculum: properties, reactivity, ionic compounds, basic metallurgy and corrosion. Practice and the native discoveries are original. The Delhi iron pillar research and industrial coating comparisons extend the core ideas; they are not extra official assessment requirements.

Physical properties describe observable behaviour without changing chemical identity. Chemical properties describe reactions. Begin with patterns, then test their limits:

• Lustre: many clean metal surfaces are shiny. A tarnished surface can hide that appearance. Iodine is a lustrous non-metal, so shine alone does not settle classification.
• Malleability and ductility: a malleable material can be shaped into thin sheets; a ductile one can be drawn into wire. Many metals show these properties, which explain foil and wire. Brittle materials tend to fracture instead.
• Conduction: metals generally conduct heat and electricity well, but their conductivities differ. Graphite conducts electricity; diamond, another form of carbon, is ordinarily an electrical insulator but an excellent heat conductor.
• State and hardness: mercury is liquid at 20°C; most familiar metals are solid. Bromine is a liquid non-metal. Sodium and potassium are soft metals, while diamond is very hard.
• Sonority: many metal objects ring when struck. The material, shape and support affect the sound, so a single silent sample is not a decisive non-metal test.

Gallium melts near 30°C, showing why a “room temperature” claim needs a temperature. This is a property to read, not a handling activity. Do not put reactive metals on skin, cut metal samples, open electrical wires or use mercury for this lesson.

Detect the mistake: “If it conducts, it must be a metal”

Graphite is a counterexample. Combining several observations is better than one test, and chemical behaviour adds further evidence. Do not replace the false rule with “all carbon conducts”: diamond and graphite have different structures and properties.

Oxygen and protective surfaces

Many metals react with oxygen to form oxides, but neither the rate nor the surface produced is identical. Copper heated in air forms black copper(II) oxide:

2Cu(s) + O₂(g) → 2CuO(s)

Magnesium burns brightly to form magnesium oxide. Aluminium forms aluminium oxide: 4Al(s) + 3O₂(g) → 2Al₂O₃(s). These are recorded classroom observations, not instructions to burn metals. A large piece and a fine powder can behave differently because their exposed surface areas differ.

Aluminium's thin, adherent oxide film helps separate the underlying metal from its environment. A metal can therefore be chemically reactive yet appear slow to react because its surface is protected. Anodising deliberately thickens the oxide layer using an electrolytic process. It is a surface treatment, not a conversion of the whole object into a less reactive element.

Metal oxides are often basic: copper oxide reacts with an acid to form a salt and water. Some are amphoteric, reacting with both acids and bases. Aluminium oxide and zinc oxide are important examples. The chapter represents aluminium oxide's reactions as:

Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O

Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O

The second is the textbook's simplified sodium-aluminate representation; species present in aqueous solution depend on conditions. “Amphoteric” describes the oxide's reactions, not a claim that aluminium metal is simultaneously an acid and a base.

Most metal oxides are poorly soluble in water. Some react with water to give soluble hydroxides, such as Na₂O + H₂O → 2NaOH. Non-metal oxides are commonly acidic or neutral; they are not all acidic and not all soluble. The oxide must be identified before predicting its behaviour.

Explain: why does a dull aluminium surface not disprove its reactivity?

The visible surface can be a protective oxide. Reactivity concerns a possible chemical change; the observed rate also depends on barriers, temperature and surface condition. The film helps explain the everyday object rather than contradicting the reactivity series.

Water and acids: conditions change the comparison

The phrase “metal reacts with water” leaves out essential information. Cold liquid water, hot water and steam give different observations. Surface coatings can also delay a reaction. The following are textbook comparisons to interpret, not experiments to reproduce.

• Sodium and potassium react vigorously with cold water, producing hydroxide and hydrogen with substantial heat. For sodium: 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g). Their reactivity explains storage under a suitable protective liquid in a laboratory.
• Calcium reacts less vigorously: Ca + 2H₂O → Ca(OH)₂ + H₂. Calcium hydroxide has limited solubility; some can remain as solid. Hydrogen bubbles adhering to the metal can help it float.
• Magnesium shows little visible reaction with cold water in the ordinary classroom comparison; hot water promotes the formation of magnesium hydroxide and hydrogen: Mg + 2H₂O → Mg(OH)₂ + H₂.
• Heated iron, zinc and aluminium can react with steam under suitable conditions. For iron: 3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g). For zinc: Zn + H₂O(g) → ZnO + H₂. For aluminium: 2Al + 3H₂O(g) → Al₂O₃ + 3H₂. A protective film matters, especially for aluminium.
• Copper, silver and gold do not liberate hydrogen from water in these ordinary comparisons. Avoid turning a classroom observation into a claim about every conceivable temperature or chemical environment.

Many metals above hydrogen in the activity series liberate hydrogen from dilute non-oxidising acids such as dilute hydrochloric acid. Zinc forms zinc chloride:

Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)

Copper does not liberate hydrogen from dilute HCl under ordinary conditions. However, “metal + any acid always gives hydrogen” is false. Nitric acid is oxidising and its reactions generally produce reduced nitrogen-containing products instead of hydrogen gas; the textbook notes magnesium and manganese with very dilute nitric acid as exceptions.

Aqua regia, mentioned in the chapter, is a highly corrosive mixture of concentrated hydrochloric and nitric acids, conventionally described in a 3:1 volume ratio. It can dissolve gold and platinum. This is a definition to recognise, not a mixing or jewellery-cleaning procedure. A shiny cleaned object can have lost metal.

Reactivity Series

The chapter's series orders common metals by decreasing activity. Hydrogen is included as a reference for acid reactions even though it is a non-metal:

K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag > Au

Connect the order to evidence. In an appropriate salt solution, a more reactive metal can displace a less reactive one. An iron nail in copper sulfate solution can acquire a copper deposit as iron(II) sulfate forms:

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

The blue solution can become paler and develop a pale-green colour as the composition changes. Copper atoms were already present as ions in solution; iron atoms did not turn into copper. Reversing the pair does not reverse the reaction: copper does not displace iron from aqueous iron(II) sulfate in this comparison.

Zn(s) + FeSO₄(aq) → ZnSO₄(aq) + Fe(s) places zinc above iron. Copper can displace silver: Cu(s) + 2AgNO₃(aq) → Cu(NO₃)₂(aq) + 2Ag(s). “Less reactive” is relative to a partner, not the same as completely inert.

The rule needs appropriate conditions. Adding sodium to an aqueous metal-salt solution involves its vigorous reaction with water, so it is not a simple test-tube route to depositing every metal below it. Protective films and observation time can also make a failed visible reaction ambiguous. Use the series with chemical context.

Build an order from evidence: zinc displaces iron; iron displaces copper

The supported order is Zn > Fe > Cu. In the usual classroom comparison, zinc should also displace copper from copper sulfate. Nothing in those two observations alone locates an untested fourth metal. That missing information is a reason to ask for another comparison, not to invent a rank.

Metals + Non Metals

Electron arrangements help explain the formation of familiar ionic compounds. Sodium has 11 protons and, when neutral, 11 electrons arranged 2,8,1. Losing one electron gives Na⁺ with 11 protons and 10 electrons, arranged 2,8. The nucleus remains sodium's nucleus.

Chlorine has 17 protons and 17 electrons arranged 2,8,7. Gaining one electron gives Cl⁻ with 18 electrons arranged 2,8,8. In electron bookkeeping, sodium supplies the electron that chlorine receives. The resulting oppositely charged ions attract.

2Na(s) + Cl₂(g) → 2NaCl(s)

The overall equation includes chlorine as Cl₂. A one-atom electron-dot picture is a way to follow valence electrons, not an alternative overall equation. Solid sodium chloride is an extended arrangement of ions, not a collection of separate NaCl molecules. NaCl gives the simplest 1:1 ratio of sodium to chloride ions.

Magnesium loses two electrons to form Mg²⁺. Two chlorine atoms each gain one to form two Cl⁻ ions. The neutral ionic formula is MgCl₂ because +2 + 2(−1) = 0. A formula such as MgCl would leave an uncompensated charge for these ions.

Make a neutral formula: combine Al³⁺ with O²⁻

Use two Al³⁺ ions and three O²⁻ ions: 2(+3) + 3(−2) = 0. The formula is Al₂O₃. For Na⁺ and O²⁻, use two sodium ions for one oxide ion: Na₂O. Charge balance determines the ratio; you do not change an ion's charge to obtain your preferred formula.

The octet model is useful for these examples, but “every atom must have eight outer electrons” is not a universal rule. Helium's filled shell contains two electrons, and more advanced bonding includes other cases. Use the model where it explains the given elements.

Properties of Ionic Compounds

Strong electrostatic attractions hold positive and negative ions in a solid arrangement. Separating them sufficiently to melt the material generally takes substantial energy, explaining the high melting temperatures of many familiar ionic solids. Ionic solids are commonly hard but brittle: shifting parts of the arrangement can bring like charges close together and favour fracture.

Having charge is not enough for electrical conduction. Charge must be able to move through the material. In ordinary solid sodium chloride the ions are held around fixed positions and cannot carry current across the solid as they can in a melt. Melting permits ion movement. Dissolving a suitable ionic compound in water also produces mobile ions.

The carriers differ from those in a metal. In a metallic wire, mobile electrons carry the current; in an ionic melt or solution, moving positive and negative ions carry it. Both signs contribute. In an applied electric field, cations move towards the negative electrode and anions towards the positive electrode.

Many ionic compounds dissolve in water, but some, including silver chloride and barium sulfate, are poorly soluble. “Ionic” does not mean “always water-soluble”. A solution's ability to conduct depends on the ions actually present and their mobility, not merely the word salt in a label.

Diagnose: a student says melting salt “creates the ions”

The ions were already present in the solid. Melting changes their mobility; it does not need to create their charges. A correct comparison identifies the same charged particles in both states and explains why they can move through one state much more readily.

Occurrence of Metals

Metals occur in naturally occurring minerals, either as elements or in compounds. An ore is material from which a valuable metal can be extracted economically under the relevant conditions. Merely containing a metal does not automatically make a rock an ore. Concentration, recovery method and cost matter.

Gold can occur in a native, uncombined form. Copper and silver occur both native and in compounds. Highly reactive metals such as sodium and aluminium occur naturally in compounds rather than as ordinary deposits of free metal. Seawater contains dissolved metal ions, including sodium and magnesium, not little pieces of those reactive metals floating in it.

Unwanted material associated with mined ore, such as sand and other rock, is called gangue. Enrichment or concentration separates useful ore material from much of the gangue using appropriate differences in properties. It is distinct from reducing metal ions to metal and from refining the resulting metal.

Follow a fictional ore batch: is concentration the same as extraction?

A 100 kg mined batch contains 20 kg of a metal-bearing mineral and 80 kg of gangue. An idealised separation removes 60 kg of gangue without losing mineral. The 40 kg concentrate is now 50% mineral rather than 20%. It is still a mineral-containing mixture, not 20 kg of pure metal. The mineral's chemical formula determines how much of its mass is metal.

Extraction: choose a route that fits reactivity

Metallurgy connects concentration, chemical extraction and refining. The activity series is a useful guide to extraction methods, although actual industrial routes depend on the ore, chemistry, energy and economics. A branching diagram summarises common routes; it should not promise one universal process for every metal in a box.

Lower activity: the mercury and copper examples

Some relatively unreactive metals can be obtained from suitable compounds by heating. The textbook describes cinnabar, HgS, first oxidising during heating in air and the oxide then decomposing:

2HgS + 3O₂ → 2HgO + 2SO₂

2HgO → 2Hg + O₂

These are chemical stages to explain, not a safe method for handling mercury: mercury vapour and the gases require industrial controls. Copper sulfide similarly illustrates partial oxidation followed by reaction between sulfide and oxide:

2Cu₂S + 3O₂ → 2Cu₂O + 2SO₂

2Cu₂O + Cu₂S → 6Cu + SO₂

Middle activity: convert to an oxide, then reduce

Roasting commonly converts a sulfide ore to an oxide by strong heating in excess air. Calcination uses heating in limited or absent air for processes such as decomposing a carbonate. The zinc examples make the contrast visible:

2ZnS + 3O₂ → 2ZnO + 2SO₂

ZnCO₃ → ZnO + CO₂

Notice oxygen as a reactant in roasting, while the carbonate already contains the oxygen in the calcination equation. The oxide can then be reduced by a suitable reducing agent. A simplified high-temperature representation is ZnO + C → Zn + CO. Carbon takes oxygen; zinc oxide loses it. The equations describe transformations, not complete furnace designs or every intermediate.

A more reactive metal can also reduce an oxide. The thermit reaction between aluminium and iron(III) oxide releases enough heat to produce molten iron and is used in specialised welding:

Fe₂O₃(s) + 2Al(s) → 2Fe(l) + Al₂O₃(s)

Aluminium is oxidised and iron(III) oxide is reduced. The chapter also gives 3MnO₂ + 4Al → 3Mn + 2Al₂O₃. Do not mix powders or attempt either reaction; our investigation concerns the equations and choice of reducing agent.

High activity: electrical energy drives reduction

For highly reactive metals, ordinary carbon-reduction routes are unsuitable. Sodium, magnesium and calcium can be obtained by electrolysis of their molten chlorides. In molten sodium chloride, Na⁺ gains an electron at the negative cathode; chloride ions lose electrons at the positive anode, producing chlorine.

2NaCl(l) → 2Na(l) + Cl₂(g)

This is an electrically driven process. Using an aqueous solution changes the chemistry because water can participate; electrolysis of brine is not a way to obtain sodium metal in water. Aluminium is obtained by electrolytic reduction of alumina in an industrial molten electrolyte. “Heat any ore with carbon” misses this important difference.

Refining: move copper, separate impurities

Extraction and purity are different questions. In electrolytic refining of copper, impure copper is the positive anode, a thin sheet of pure copper is the negative cathode, and acidified copper sulfate solution is the electrolyte.

At the anode, copper atoms lose electrons and enter solution as Cu²⁺ ions. At the cathode, Cu²⁺ ions gain electrons and deposit as copper metal. The anode becomes smaller and the cathode gains copper. Electrons pass through the external circuit; ions carry charge through the electrolyte.

Impurities behave according to their chemistry: some dissolve and remain in solution under the refining conditions, while others collect beneath the anode as anode mud. The process does not deposit every dissolved metal indiscriminately. It is controlled to obtain purer copper.

Predict: in an ideal copper-only transfer, the anode loses 5 g of copper

The cathode gains 5 g of copper if current efficiency is ideal and there are no side reactions or other copper losses. That does not imply that an impure anode's total mass loss equals the copper deposited: separately account for impurity material. The model tracks copper, not every component as though it were copper.

Corrosion and prevention

Corrosion is a metal's deterioration through reactions with its environment. Ordinary iron rusting needs oxygen and moisture. In the textbook's comparison, iron with both air and water rusts; dry air and water with dissolved oxygen largely removed give much less rust under the observation conditions.

The comparison works by controlling factors. The absence of rapid visible rust over a short observation is not proof of zero reaction forever. Rust's porous, flaky nature allows further attack, unlike an effective protective film. Other metals form different corrosion products: silver tarnish commonly contains silver sulfide; copper's green surface can contain basic copper carbonate and other compounds depending on the environment.

• Paint, oil or grease: a suitable maintained barrier reduces exposure. Scratching or loss of coverage can expose metal.
• Galvanising: a zinc coating provides a barrier and can sacrifice zinc to protect small exposed regions of steel. Protection after damage is limited by the remaining zinc, size of damage and environment.
• Anodising: a thicker protective oxide improves aluminium's surface resistance.
• Alloy selection: suitable stainless steels resist many environments through a chromium-rich passive film. The grade and service conditions matter.

Stainless does not mean incapable of corrosion. Worldstainless describes conditions that can cause pitting, crevice attack and other corrosion. That caveat makes material selection more useful: a coastal fitting and an indoor object need not face the same exposure.

Alloys: properties can be designed

An alloy is a metallic material made by combining a metal with one or more other elements. The school description treats alloys as mixtures of metals, or of a metal and a non-metal. Unlike an ionic compound such as MgCl₂, an alloy name generally does not specify one fixed small-number atomic formula.

Brass combines copper and zinc; traditional bronze combines copper and tin. Steel is based on iron with carbon, often with other additions. Stainless steels contain chromium, and some grades also contain nickel or other elements. Nickel is not a compulsory ingredient of every stainless steel. An alloy containing mercury is called an amalgam. Traditional tin–lead solder illustrates a low-melting alloy used for joining metals; this is a material example, not an instruction to handle or heat lead.

Alloying can change strength, hardness, corrosion resistance, conductivity and melting behaviour. The direction and size of a change depend on composition and processing. Brass still conducts electricity, although less well than pure copper; calling it a non-conductor would be misleading. Some solders melt at useful low temperatures, but not every alloy has a lower melting temperature than each constituent metal.

Read the proportion: what does 22-carat gold mean in an ideal alloy calculation?

It indicates 22 parts gold out of 24 parts by mass: 22 ÷ 24 × 100 ≈ 91.7%. An ideal 12 g sample therefore contains 11 g gold and 1 g other alloying material. This is a proportion exercise, not a method to authenticate a jewellery item.

Discover it: what does a protective layer explain?

Extension: the Delhi iron pillar. Its exceptional atmospheric corrosion resistance is not evidence that iron there cannot react. Research by IIT Kanpur metallurgist R. Balasubramaniam links its resistance to the development of a protective surface film, with the iron's phosphorus content and atmospheric exposure playing important roles.

Compare that with modern stainless steel's chromium-rich film and zinc's sacrificial protection. All three concern a surface, but they are not the same mechanism or composition. The comparison connects Indian metallurgical history with a general engineering question: which material and environmental conditions make a protective system work?

In the accompanying native workshop, arrange virtual reactivity evidence, move paper ions between fixed and mobile states, and challenge a corrosion claim. Write one observation that would distinguish two explanations. The optional original electron-dot diagram shows a useful bookkeeping convention; explain why the ionic solid still needs an extended arrangement.

Final transfer: a coating looks intact. Does that guarantee protection?

No. Ask what the coating is, how it works, whether the environment permits that mechanism and whether damage or wear has changed it. Appearance is evidence, but material identity, exposure and inspection matter too. Do not test a household coating with acids or scratch a real object; build the argument on paper.

• Physical properties are patterns with exceptions; particle structure and surfaces explain why.
• Use the activity series with appropriate reagents and conditions.
• Ionic formulae balance charge; conduction requires mobile charge carriers.
• Concentration, extraction and refining solve different problems.
• Corrosion protection is a mechanism under conditions, not a permanent guarantee from a name.

Sources

Core scope and chemistry: NCERT, Metals and Non-metals, Reprint 2026–27; CBSE Science curriculum, 2026–27. Revisit Chemical Reactions and Equations, Acids, Bases and Salts or Electricity for the shared concepts.

Property and source checks: Royal Society of Chemistry on carbon's forms and properties and gallium; OpenStax solubility and redox explanations; Worldstainless on stainless steel families and corrosion conditions; American Galvanizers Association on damaged zinc coatings.

Additional checks: Princeton on copper–nickel phase behaviour; Canadian Conservation Institute on silver tarnish (the linked conservation activity is not part of this paper-only lesson).

Historical science extension: Balasubramaniam, protective film formation on the Delhi iron pillar, Current Science, 2003. Return to the chemistry topic hub or study library.

Test yourself

What is the main difference between metals and non-metals?

The main difference between metals and non-metals lies in their physical and chemical properties, such as lustre, malleability, ductility, conduction, and reactivity.

Why does a dull aluminium surface not disprove its reactivity?

A dull aluminium surface does not disprove its reactivity because the visible surface can be a protective oxide film that separates the underlying metal from its environment.

What is the significance of the oxide film on aluminium?

The oxide film on aluminium helps to protect the metal from further corrosion by separating it from its environment.

How do metal oxides react with acids and bases?

Metal oxides can be basic, amphoteric, or neutral, and their reactions with acids and bases vary accordingly.

Why is it important to consider multiple properties when classifying metals and non-metals?

Considering multiple properties is essential for accurate classification because a single property, such as conduction, can be misleading, as seen in the case of graphite and diamond.