Chapter 6 — Metals and Non-Metals

This topic introduces the two broad groups of elements commonly encountered in everyday life: metals and non-metals. We start by asking simple observational questions: does the substance shine when polished? Can it be hammered into a thin sheet or drawn into a wire? Does it conduct heat and electricity? The answers to such questions help us place an element into the metal or non-metal group. Metals are often shiny, malleable and good conductors; non-metals are more likely to be dull, brittle and poor conductors. Some non-metals are gases (oxygen, nitrogen), one is a liquid at room temperature (bromine), and many are solids with varied appearances (sulphur, carbon).

It is useful to remember everyday examples: iron, copper, aluminium and gold are common metals; carbon (as charcoal), sulphur, oxygen and chlorine are common non-metals. While the periodic table shows where metals and non-metals lie, at Class 7 the emphasis is on recognising and using properties to sort materials. Classification is not only an academic exercise — it explains why cooking utensils are made of metals, why electrical wires use copper, and why rubber or plastic (non-metals) are used as handles or insulators. The topic also sets the stage for later ideas: metals often form basic oxides and ionic salts with non-metals, and their different behaviours determine extraction methods, protection from corrosion and practical applications.

Practical skills developed here include observing physical properties, recording results, and making simple predictions. These skills are important for carrying out safe classroom experiments and for understanding real-world uses of materials. Always perform tests carefully and under teacher supervision, especially when using heat, acids or reactive metals.

This topic explores in detail the physical properties that most metals share, and how to test them safely in the classroom. The first property is lustre: most metals look shiny when freshly cut or polished. This shininess results from the way the metal surface reflects light. Second, malleability means a metal can be beaten into thin sheets without breaking; gold and aluminium are typical malleable metals. Third, ductility describes the ability to be drawn into wires — copper is widely used in electrical wiring for this reason. Conductivity is another key feature: metals are generally good conductors of heat and electricity, which is why metal pots conduct heat quickly and metal wires carry electric current.

Other common properties include hardness and density, though these can vary widely. Some metals like sodium and potassium are soft and can be cut with a knife; others like iron and tungsten are hard. Most metals are solids at room temperature; mercury is a notable liquid metal. Colour varies too: most are silvery, but copper is reddish and gold is yellow. When testing properties, simple experiments include: polishing a sample to check lustre, tapping it to see if it bends (malleability) or breaks, and connecting it in a simple electrical circuit to see if it lights a bulb (conductivity).

Understanding these properties explains everyday choices. Malleability allows metals to be formed into sheets for cooking vessels; ductility allows wires to be made for homes and industries; conductivity determines which metals are used in electrical work. Note exceptions and safety: some metals react with air or water and should be handled by teachers during demonstrations. Always record observations carefully and compare results for different metals to see patterns and exceptions.

Non-metals show a wide range of physical properties that contrast with metals. Many non-metals are dull — they do not have the shiny appearance that metals have. In solid form they are often brittle, meaning they break or crumble rather than bend. For example, a piece of sulphur powder or phosphorus will break rather than deform. Non-metals are usually poor conductors of heat and electricity; this makes them useful in insulating materials — substances like rubber and plastic prevent the flow of electricity and protect us from shocks.

Non-metals can be gases, liquids or solids at room temperature. Oxygen and nitrogen are gases essential for life and industry; bromine is a red-brown liquid at room temperature; carbon (in forms like charcoal or graphite) and sulphur are solids. Their melting and boiling points vary widely: many non-metallic gases have low boiling points, while diamond (a form of carbon) has an extremely high melting point and is very hard. The appearance and texture also differ — carbon (charcoal) is black and powdery, iodine sublimes to violet vapour, and sulphur is bright yellow.

Classroom tests for non-metals include trying to conduct electricity, checking brittleness by careful tapping, and observing state at room temperature. Many non-metals do not react with water or acids the way metals do. Understanding these varied properties helps decide where non-metals are used: gases for breathing and combustion, solids for insulation or chemical reagents, and liquids like bromine for industrial chemistry (handled only by teachers). Note that some non-metals like graphite conduct electricity due to their special structure, and such exceptions are important to remember.

When metals react with oxygen, they form compounds called oxides. This chemical change is known as oxidation. Some reactions with oxygen are rapid and produce flames and bright light, while others are slow and may only produce a coloured layer on the metal's surface. A classic dramatic classroom example is burning magnesium ribbon: it burns with a bright white flame and leaves a white powder that is magnesium oxide.

Other metals like iron react slowly with oxygen, often in the presence of water, to form rust (a hydrated iron oxide). Copper when heated in air forms copper oxide which is black; further exposure can give green copper carbonate in the presence of carbon dioxide and moisture. The products — metal oxides — have different properties from the original metal: many metal oxides are basic in nature. Observational clues to oxidation include change in colour, formation of powder or crust, release of heat and sometimes light.

At Class 7 level, students learn word equations for common oxidation reactions and practise writing them. They also learn that reactivity varies: very reactive metals like potassium and sodium oxidise quickly and violently; less reactive ones like gold do not oxidise easily and are often found as native metals in nature. Experiments should be done by the teacher or under strict supervision because some oxidations are energetic. Finally, understanding oxidation explains everyday phenomena such as tarnishing of metals, formation of patina on statues, and the need for protective coatings on metal surfaces.

Metals show different behaviour when they meet water, depending on how reactive they are. Very reactive metals such as potassium, sodium and calcium react vigorously with cold water to form a metal hydroxide and hydrogen gas. These reactions may be explosive and are therefore always demonstrated by the teacher with safety precautions. Less reactive metals like iron react slowly with water, especially when hot, and some metals like copper do not react with water under normal conditions.

Typical classroom observations for reactions with water include bubbling (release of hydrogen gas), change in temperature (often warming due to exothermic reaction), and sometimes a change in the appearance of the metal as it dissolves or forms a coating. For sodium, the reaction with water is vigorous: sodium floats, moves on the water surface, produces bubbles and may ignite the hydrogen produced. Calcium reacts less violently but still produces hydrogen and forms calcium hydroxide which makes the water slightly alkaline.

Word equations help express these reactions: for example sodium + water → sodium hydroxide + hydrogen. Students should practise writing such word equations and recognise that the products often include a hydroxide (metal + water → metal hydroxide + hydrogen) for reactive metals. The topic also links to safety and practical consequences: metals that react with water are unsuitable for plumbing or storage in wet places unless protected, and certain metals are stored under oil to avoid contact with moisture in air.

Reactions of metals with dilute acids are a standard test to compare metal reactivity and to observe gas evolution. When a reactive metal is dropped into dilute acid, the usual products are a salt and hydrogen gas. For example, zinc reacts with dilute hydrochloric acid to give zinc chloride and hydrogen; magnesium reacts with dilute sulfuric acid to form magnesium sulfate and hydrogen. These reactions can be represented by simple word equations, and they show variations in reaction rate depending on the metal.

Classroom observations include bubbling as hydrogen gas is produced, warming of the solution in many cases (exothermic reaction), and gradual disappearance of the metal if it dissolves. The rate of reaction relates to reactivity: magnesium reacts quickly, zinc moderately, and iron more slowly with dilute acids; copper, silver and gold generally do not react with dilute acids. The gas produced can be tested by collecting it over water or by bringing a lit splint near the mouth of the test-tube; a characteristic 'pop' sound confirms hydrogen. Always follow strict safety rules: use dilute acids only under the teacher's supervision, wear safety goggles and gloves, and keep flames away from the reaction if hydrogen is evolved.

It is helpful to know how salts are named in these reactions: hydrochloric acid produces chloride salts (e.g., zinc chloride), and sulfuric acid produces sulfate salts (e.g., magnesium sulfate). These salts often dissolve in water, changing the colour or conductivity of the solution. Learning these acid-metal reactions helps students place metals in a simple reactivity order, understand industrial uses of acids in metal cleaning or processing, and practise writing and balancing word equations. Teachers should demonstrate any vigorous reactions and discuss real-life implications such as acid damage to metal structures and the need for protective measures.

Displacement reactions reveal how metals differ in their tendency to lose electrons and form ions. In a typical displacement reaction, a more reactive metal will displace a less reactive metal from a solution of its salt. For example, when an iron nail is placed in copper sulfate solution, iron displaces copper and a brown deposit of copper forms on the nail. This observation helps students understand that iron is more reactive than copper.

A reactivity series is a simple vertical list of metals arranged from most reactive to least reactive. Although a full, precise series is taught in higher classes, Class 7 students learn a basic order: very reactive metals (e.g., potassium, sodium, calcium) at the top, moderately reactive metals (e.g., magnesium, aluminium, zinc, iron) in the middle, and less reactive metals (e.g., copper, silver, gold) at the bottom. The series helps predict whether a displacement reaction will occur: a metal can displace ions of another metal that is below it in the series.

Experiments with displacement reactions are simple and visual: placing strips of different metals in solutions of metal salts shows colour changes and deposition. These activities teach the idea of chemical reactivity in a practical way and link to extraction of metals (more reactive metals are harder to extract) and corrosion (less reactive metals resist corrosion better). Students should practise writing word equations for displacement reactions and use the reactivity series to predict results.

This topic explains how metals and non-metals react together to form ionic compounds, a common type of chemical substance. In such reactions, metal atoms tend to lose electrons and form positive ions (cations), while non-metal atoms gain electrons and form negative ions (anions). The electrostatic attraction between oppositely charged ions holds them together, producing a stable ionic lattice. A well-known everyday example is sodium reacting with chlorine to form sodium chloride (table salt).

Properties of ionic compounds are characteristic and different from pure metals: they often form hard, crystalline solids with definite shapes and high melting and boiling points because the ionic bonds are strong. In the solid state ionic compounds do not conduct electricity because the ions are fixed in place; however, when melted or dissolved in water the ions are free to move and the compound conducts electricity. This behaviour is important for understanding why salts conduct only in solution and why metals conduct in the solid state due to mobile electrons.

For Class 7, students learn to write simple word equations such as sodium + chlorine → sodium chloride and to recognise common ionic products like metal oxides and salts. Drawing a simple diagram helps: show a metal atom losing an electron and the non-metal gaining it, label the resulting cation and anion and indicate the attraction between them. Discuss practical consequences: many ores are metal compounds which are converted into ionic salts during processing; ionic salts are used in food, medicine and industry. Teachers should explain that ionic bonding is a strong force and that ionic compounds are typically brittle — a hammer blow can shatter the crystal because layers of ions with like charges may be forced together and repel each other.

Metals are not evenly distributed in the Earth's crust. They occur either as native metals — found in pure or nearly pure metallic form — or as compounds in minerals and ores. Native metals such as gold and sometimes silver are less reactive and can be found as nuggets or grains in riverbeds and rocks. Most metals are found combined with other elements in minerals; when a mineral contains enough of the metal to make mining worthwhile, we call it an ore.

Different metals occur in different types of ores. For example, iron is commonly found in ores such as hematite and magnetite; aluminium is obtained from bauxite which contains aluminium oxide compounds; copper occurs in ores like chalcopyrite. The form in which a metal occurs affects how easily it can be extracted: metals close to the bottom of the reactivity series (less reactive) are often easier to extract because they can be reduced by heating, while reactive metals may need special chemical or electrical methods.

Understanding occurrence links chemistry to geography and industry. Mining extracts ores from the ground; ores must be concentrated and then processed to obtain the metal. Mining affects the environment through landscape change and waste, so responsible extraction with reclamation and pollution control is important. For Class 7, remember common ores and the idea that ores must be processed to get the metal whereas native metals may sometimes be used with little or no refining.

This topic gives a simple overview of how metals are obtained from their ores. The method used depends largely on the metal's reactivity. Less reactive metals like iron and copper can often be obtained by heating their ores with carbon (a reduction process); carbon removes oxygen from the metal oxide to leave the metal. In contrast, very reactive metals such as aluminium cannot be extracted by carbon reduction and require electrolysis, a process that uses electricity to separate the metal from its compound.

There are several general steps in metal extraction: concentration of the ore to remove impurities, conversion of the ore to a suitable form (such as an oxide), reduction to the metal (by heating with carbon or by electrolysis), and purification of the metal. For example, iron is extracted from its ore in a blast furnace where coke (carbon) reduces iron oxide to iron. Aluminium is obtained by electrolysis of molten alumina because aluminium is too reactive to be reduced by carbon.

Energy use and environmental impact are important practical concerns: extraction often requires high temperatures or large amounts of electricity and produces wastes. Recycling metals reduces the need for mining and saves energy. For Class 7, the aim is to understand the basic ideas (concentration, reduction, electrolysis) and to be able to give simple examples of how common metals like iron and aluminium are obtained and why the method differs depending on reactivity.

Corrosion means the gradual chemical attack or destruction of a metal by substances in its surroundings. The most common example is rusting of iron, which occurs when iron reacts with oxygen and water to form hydrated iron(III) oxide — a flaky, reddish-brown deposit that weakens the metal. Corrosion can affect many metals and can take several forms: tarnishing of silver, formation of a green patina on copper, or pitting corrosion on metals exposed to salty air.

The factors that increase corrosion include moisture, oxygen, salts (like seawater), acids, and high temperatures. Salt speeds up corrosion because dissolved salts increase the electrical conductivity of water and so promote electrochemical reactions on the metal surface. Acidic conditions accelerate attack because acids provide more hydrogen ions which participate in corrosion processes. Mechanical damage to protective coatings and scratches expose fresh metal and can start rapid corrosion.

Several methods are used to prevent or slow corrosion. These include coating the metal with paint or oil to keep out air and water; galvanising, which covers iron with a layer of zinc that corrodes preferentially and thus protects the iron; and alloying, for example adding chromium to make stainless steel which forms a thin protective oxide layer that prevents further corrosion. Cathodic protection uses a more reactive metal connected to the metal to be protected, so the sacrificial metal corrodes instead. Regular maintenance, drying, and keeping metals away from salt or acids are practical steps used at home and in industry. Classroom demonstrations — such as putting identical iron nails in different environments (air, water, salt water, oil) — help students see how conditions change the rate of rusting. Understanding corrosion is important for safety and economics: prevention saves money and prevents accidents caused by weakened structures.

An alloy is a mixture of a metal with one or more other elements to improve its properties. Alloys are designed to be stronger, harder, less likely to corrode, or better suited for particular uses than pure metals. For example, steel is an alloy of iron with a small percentage of carbon which makes it harder and stronger; it is used in buildings, tools and vehicles. Brass (copper and zinc) is harder than copper and has a pleasing colour, so it is used for decorative fittings and musical instruments. Bronze (copper and tin) is durable and used for statues and bearings.

Knowing the properties of metals and non-metals explains their everyday uses. Metals are chosen where strength, conductivity or malleability is needed: aluminium for light cooking utensils and aircraft parts because it is light and conducts heat well; copper for electrical wiring because it is an excellent conductor and ductile; gold and silver for jewellery because they resist corrosion and have attractive appearance. Non-metals are chosen where insulation or chemical behaviour is required: rubber and plastic for insulating handles and electric cables, oxygen for respiration and combustion, nitrogen for making fertilizers, and sulphur for matches and chemicals.

The right material is chosen considering cost, availability and resistance to wear or corrosion. Recycling metals saves energy and reduces environmental damage from mining. For Class 7, be ready to match properties to uses and to name common alloys and products made from metals and non-metals. Practical decisions such as choosing copper for wiring or plastics for insulation follow directly from the properties studied in this unit.