Metals and Non-metals

Metals

Have you ever wondered what makes the gold in jewelry shine, the copper in our wires conduct electricity, or the iron in a bridge so strong? These materials, and many more like them, belong to a vast and vital family of elements called metals.

In our world, elements are broadly sorted into two main groups: metals and non-metals. This chapter is your guide to understanding these fundamental building blocks of matter. We will explore their distinct characteristics, how they interact, and how we extract them from the earth to build our modern world.

{{VISUAL: photo: a collage of everyday objects made from different metals (gold ring, copper wire, aluminum foil) and non-metals (graphite pencil lead, sulfur powder, diamond)}}

Throughout this chapter, we will uncover:

• The physical and chemical properties that set metals and non-metals apart.
• How the reactivity series helps us predict chemical reactions.
• The formation of ionic compounds through the transfer of electrons.
• The fascinating process of metallurgy: extracting metals from their ores.
• The real-world problem of corrosion and how to prevent it.

{{KEY: type=definition | title=Metals | text=Metals are elements that are typically hard, shiny, malleable, fusible, and ductile, with good electrical and thermal conductivity.}}

Understanding these properties is not just about chemistry; it's about seeing why we choose specific materials for specific jobs—from a sturdy steel beam to a delicate aluminum foil. The unique features of metals make them indispensable to our daily lives and technological progress.

{{VISUAL: chart: simple bar chart comparing the electrical conductivity of copper (metal), silicon (metalloid), and sulfur (non-metal) to show the vast difference}}

Now, let's begin our journey by taking a closer look at the physical properties that make metals so uniquely useful.

Non-metals

Page 2: The World of Non-metals

On the previous page, we explored the shiny, strong, and conductive world of metals. Now, let's turn our attention to their counterparts: the non-metals. If metals are the sturdy pillars of the periodic table, non-metals are the versatile and essential gases, liquids, and brittle solids that make life possible. Think of the oxygen you breathe, the carbon that forms the backbone of all life, or the chlorine that keeps our water safe. These are all non-metals!

Located primarily on the upper right side of the periodic table, non-metals have properties that are often the complete opposite of metals.

Physical Properties of Non-metals

While metals share a fairly consistent set of physical traits, non-metals are much more diverse. They can be gases, a liquid, or solids at room temperature. Let's break down their typical characteristics.

• Physical State: Non-metals exist in all three states. For example, oxygen and nitrogen are gases, bromine is a liquid, and carbon, sulphur, and phosphorus are solids.
• Lustre: They are typically dull and do not have a shiny surface. The one major exception is iodine, which has a lustrous, crystalline appearance.
• Hardness: Non-metals are generally soft. However, there's a huge exception: diamond, an allotrope (different form) of carbon, which is the hardest natural substance known.
• Malleability and Ductility: Non-metals are brittle, meaning they break or shatter when hammered or stretched. You can't make a sheet or a wire out of sulphur or solid carbon.
• Conductivity: They are poor conductors of both heat and electricity. But again, we have a critical exception: graphite, another allotrope of carbon, is an excellent conductor of electricity and is used to make electrodes.
• Density: They generally have low densities and are light for their size.
• Sonority: Non-metals are not sonorous. They do not produce a ringing sound when struck.
• Melting and Boiling Points: Most non-metals have relatively low melting and boiling points compared to metals. (Diamond and graphite are notable exceptions with very high melting points).

{{VISUAL: photo: a collage showing different non-metals in their natural state - yellow sulphur powder, black carbon (coal), reddish-brown bromine liquid in a sealed ampoule, and a representation of invisible oxygen gas.}}

{{KEY: exam | title=Exceptions are Key Questions | text=CBSE frequently asks questions based on the exceptions to the general properties. Memorize these: Iodine (lustre), Diamond (hardness), and Graphite (electrical conductivity).}}

{{ZOOM: title=Allotropes of Carbon | text=Why are diamond and graphite so different, even though both are just carbon? It's all about how the atoms are arranged. In diamond, each carbon atom is tightly bonded to four others in a rigid 3D structure, making it extremely hard. In graphite, carbon atoms are arranged in flat hexagonal sheets that can slide over each other, making it soft and a good conductor because of free-moving electrons within the layers.}}

Chemical Properties of Non-metals

The chemical behaviour of non-metals is dominated by their tendency to gain or share electrons to form stable electron configurations.

1. Reaction with Oxygen

When non-metals burn in oxygen, they form non-metallic oxides. These oxides are typically acidic or neutral in nature.

• Acidic Oxides: When these oxides dissolve in water, they form acids. For example, sulphur burns in air to produce sulphur dioxide gas. This gas dissolves in water to form sulphurous acid, which will turn blue litmus paper red. S (s) + O₂ (g) → SO₂ (g) SO₂ (g) + H₂O (l) → H₂SO₃ (aq) (Sulphurous Acid)

• Neutral Oxides: Some non-metal oxides do not show any acidic or basic properties. They are neutral. Common examples include carbon monoxide (CO), nitrous oxide (N₂O), and water (H₂O) itself.

{{VISUAL: diagram: a step-by-step diagram showing how to test for the acidic nature of a non-metal oxide. Step 1: Burning sulphur powder on a deflagrating spoon in air. Step 2: Placing the spoon inside a gas jar to collect the sulphur dioxide (SO₂) gas. Step 3: Adding a small amount of water and shaking to dissolve the gas. Step 4: Dipping a strip of blue litmus paper into the solution, showing it turn red.}}

{{KEY: concept | title=Nature of Non-metal Oxides | text=Non-metals react with oxygen to form acidic or neutral oxides. Acidic oxides dissolve in water to form acids. This is a key difference from metallic oxides, which are typically basic.}}

2. Reaction with Water and Dilute Acids

Generally, non-metals do not react with water or dilute acids to produce hydrogen gas.

Why? To displace hydrogen from water or acids, an element needs to be able to give electrons to the hydrogen ions (H⁺). Non-metals are electron acceptors (they are electronegative), not electron donors. Therefore, they cannot supply the electrons needed to reduce the hydrogen ions to hydrogen gas.

3. Reaction with Salt Solutions

Just like with metals, a more reactive non-metal can displace a less reactive non-metal from its salt solution. This is most commonly seen with the halogens (Group 17 elements).

For example, chlorine is more reactive than bromine. So, if you bubble chlorine gas through a solution of sodium bromide, the chlorine will displace the bromine.

Cl₂ (g) + 2NaBr (aq) → 2NaCl (aq) + Br₂ (aq) (Greenish-yellow gas) + (Colourless solution) → (Colourless solution) + (Reddish-brown liquid)

{{KEY: points | title=Properties of Non-metals | text=- Typically exist as solids, liquids or gases.

• Are brittle, non-lustrous, and poor conductors (with key exceptions).
• Form acidic or neutral oxides with oxygen.
• Do not react with water or dilute acids to displace hydrogen.
• More reactive non-metals can displace less reactive ones from their salt solutions.}}

4. Reaction with Hydrogen and Chlorine

Non-metals react with hydrogen and other non-metals like chlorine to form covalent compounds by sharing electrons.

• With Hydrogen: They form covalent hydrides. H₂ (g) + S (l) → H₂S (g) (Hydrogen Sulphide)

• With Chlorine: They form covalent chlorides. P₄ (s) + 6Cl₂ (g) → 4PCl₃ (l) (Phosphorus trichloride)

Non-metals complete their story not by giving away what they have, but by gaining or sharing to build the molecules essential for the world around us.

What happens when Metals are burnt in Air?

{{FORMULA: expr=2Cu(s) + O₂(g) → 2CuO(s) | symbols=Cu:Copper, O₂:Oxygen, CuO:Copper(II) Oxide}}

What Happens When Metals are Burnt in Air?

Have you ever seen a blacksmith heating a piece of iron until it glows red? Or perhaps you've noticed how a silver spoon tarnishes over time? These are everyday examples of metals interacting with the air around us, specifically with oxygen. Let's dive deep into this fundamental chemical property.

When most metals are burnt in air, they combine with oxygen to form metal oxides. The general reaction can be written as:

Metal + Oxygen → Metal Oxide

This might seem simple, but the real story is in the details. The speed and nature of this reaction vary dramatically from one metal to another.

Let's start with a classic school laboratory experiment. If you take a strip of magnesium ribbon and heat it over a flame, it burns with a dazzling, brilliant white light, leaving behind a white powdery ash. This ash is magnesium oxide (MgO).

2Mg(s) + O₂(g) → 2MgO(s)

{{VISUAL: photo: A magnesium ribbon being held by tongs and burning with a dazzling white flame against a dark background, producing white smoke and ash.}}

The Nature of Metal Oxides: Basic or Acidic?

So, we've formed a metal oxide. What is its chemical nature? Let's find out. If you take the magnesium oxide ash and dissolve it in water, it forms magnesium hydroxide (Mg(OH)₂), which is a base.

MgO(s) + H₂O(l) → Mg(OH)₂(aq)

How can we confirm it's a base? Simple! If you test this solution with litmus paper, you will find that it turns red litmus paper blue, a classic sign of a basic substance.

{{KEY: type=definition | title=Metal Oxides | text=Compounds formed when a metal reacts with oxygen. Most metal oxides are basic in nature because they react with water to form bases.}}

Most metal oxides are basic oxides. However, some metals, like sodium and potassium, form oxides that are highly soluble in water. These solutions are called alkalis.

• Sodium Oxide: Na₂O(s) + H₂O(l) → 2NaOH(aq) (Sodium Hydroxide)
• Potassium Oxide: K₂O(s) + H₂O(l) → 2KOH(aq) (Potassium Hydroxide)

The Curious Case of Amphoteric Oxides

Chemistry loves its exceptions! While most metal oxides are basic, some exhibit a fascinating dual nature. They can react with both acids and bases to produce salt and water. Such metal oxides are called amphoteric oxides.

The two most important examples for you to remember are aluminium oxide (Al₂O₃) and zinc oxide (ZnO).

Let's see how aluminium oxide behaves:

1. Reaction with an Acid (HCl): Here, Al₂O₃ acts like a base. Al₂O₃(s) + 6HCl(aq) → 2AlCl₃(aq) + 3H₂O(l) (Aluminium Oxide + Hydrochloric Acid → Aluminium Chloride + Water)

2. Reaction with a Base (NaOH): Here, Al₂O₃ acts like an acid. Al₂O₃(s) + 2NaOH(aq) → 2NaAlO₂(aq) + H₂O(l) (Aluminium Oxide + Sodium Hydroxide → Sodium Aluminate + Water)

This dual behaviour is a crucial property and a favourite topic for board exam questions.

{{VISUAL: diagram: Two beakers side-by-side. In beaker 1, aluminium oxide powder is shown reacting with HCl to form AlCl₃ and water. In beaker 2, the same aluminium oxide powder is shown reacting with NaOH to form sodium aluminate and water. Both beakers are clearly labeled to show the dual acidic and basic reactions.}}

{{KEY: type=concept | title=Amphoteric Oxides | text=Metal oxides that react with both acids and bases to produce salt and water are known as amphoteric oxides. They show both acidic and basic properties. Key examples are Aluminium Oxide (Al₂O₃) and Zinc Oxide (ZnO).}}

Do All Metals React with Oxygen in the Same Way?

Not at all! The reactivity of metals towards oxygen varies widely. This difference is a key indicator of their position in the reactivity series.

• Vigorous Reaction: Metals like Potassium (K) and Sodium (Na) are so reactive that they react violently with oxygen at room temperature and catch fire if left in the open. To prevent this, they are stored under kerosene oil.

• Protective Layer Formation: Metals like Magnesium (Mg), Aluminium (Al), Zinc (Zn), and Lead (Pb) are also highly reactive. However, when they react with oxygen, they form a very thin, tough, and non-porous layer of oxide on their surface. This layer acts like a protective shield, preventing the metal underneath from further oxidation or corrosion. This is why an aluminium utensil doesn't just "disappear" over time despite being a reactive metal.

• Reaction on Heating: Iron (Fe) does not burn in air even on strong heating, but iron filings burn vigorously when sprinkled into a flame. Copper (Cu) does not burn but gets coated with a black layer of copper(II) oxide (CuO) when heated for a long time. 2Cu + O₂ → 2CuO

• No Reaction: Metals like Silver (Ag) and Gold (Au) are at the bottom of the reactivity series. They do not react with oxygen even at very high temperatures, which is why they are used to make jewellery and are considered "noble" metals.

{{KEY: type=points | title=Reactivity of Metals with Air | text=- Potassium & Sodium: React violently, catch fire, stored in kerosene.

• Magnesium, Aluminium, Zinc, Lead: Form a thin, protective oxide layer that prevents further corrosion.
• Iron: Does not burn as a block, but iron filings burn in a flame.
• Copper: Forms a black coating of Copper(II) Oxide on heating.
• Silver & Gold: Do not react with oxygen.}}

This varying reactivity with oxygen is one of the most important clues that helped scientists arrange metals into the reactivity series. A more reactive metal will react with oxygen more readily and vigorously.

What happens when Metals react with Water?

{{FORMULA: expr=Metal + Water → Metal Hydroxide + Hydrogen | symbols=Metal:a reactive metal, Water:H₂O, Metal Hydroxide:a base formed by the metal, Hydrogen:H₂ gas}}

What Happens When Metals React with Water?

Imagine dropping a piece of metal into a glass of water. What do you expect to see? A violent explosion? A gentle fizz? Or nothing at all? The answer, interestingly, is all of the above! The way a metal reacts with water is a direct indicator of its reactivity. This reaction is a classic example of a single displacement reaction where a more reactive metal displaces hydrogen from water.

The general outcome is the formation of a metal hydroxide (if the hydroxide is soluble) or a metal oxide, along with the liberation of hydrogen gas (H₂). Let's explore how different metals behave, from the most dramatic to the completely inert.

1. The Explosive Ones: Reaction with Cold Water

The most reactive metals, like Potassium (K) and Sodium (Na), are at the top of the reactivity series. They react vigorously, even violently, with cold water.

The reaction is so intensely exothermic (releases a large amount of heat) that the hydrogen gas produced immediately catches fire.

• Sodium (Na): When a small piece of sodium is dropped in water, it darts around the surface, melts into a silvery globule due to the heat, and fizzes rapidly, producing hydrogen gas. This gas often ignites with a golden-yellow flame. 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) + Heat energy

• Potassium (K): Potassium's reaction is even more violent than sodium's. The hydrogen gas produced catches fire almost instantly, burning with a lilac (pale violet) flame. 2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g) + Heat energy

{{VISUAL: photo: a small piece of sodium metal reacting vigorously in a beaker of water, producing fizzing bubbles and a small orange flame on the surface.}}

2. The Fizzers: A Calmer Reaction with Cold Water

Moving down the reactivity series, we find metals that react with cold water, but much less violently. Calcium (Ca) is a prime example.

When calcium is added to water, it reacts steadily to form calcium hydroxide and hydrogen gas. The reaction is less exothermic, so the heat produced is not enough to ignite the hydrogen gas.

Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)

An interesting observation here is that the piece of calcium metal starts to float on the surface of the water after some time. Why do you think this happens?

{{ZOOM: title=Why do Calcium and Magnesium Float? | text=The bubbles of hydrogen gas produced during the reaction stick to the surface of the metal. These gas bubbles increase the buoyancy of the metal piece, making it lighter than the volume of water it displaces, causing it to rise and float on the surface.}}

{{KEY: points | title=Observations for Calcium Reaction | text=- Bubbles of a colourless gas are seen.

• The solution turns milky or cloudy due to the formation of sparingly soluble calcium hydroxide (slaked lime).
• The piece of calcium metal begins to float.}}

3. The Hot Water Enthusiasts

Some metals are a bit picky and prefer their water warm. Magnesium (Mg) does not react with cold water but reacts readily with hot water.

Similar to calcium, it forms magnesium hydroxide and hydrogen gas. It also starts to float for the same reason—hydrogen bubbles sticking to its surface.

Mg(s) + 2H₂O(l) (Hot) → Mg(OH)₂(aq) + H₂(g)

4. The Steam-Powered Reactions

Metals like Aluminium (Al), Zinc (Zn), and Iron (Fe) are even less reactive. They do not react with cold or hot water. They require the high energy of steam (gaseous water, H₂O(g)) to react.

A crucial difference here is the product formed. Instead of forming metal hydroxides, these metals form metal oxides and hydrogen gas. Metal hydroxides are generally unstable at the high temperatures required for the reaction with steam and decompose into oxides.

• Aluminium: 2Al(s) + 3H₂O(g) → Al₂O₃(s) + 3H₂(g)
• Iron: 3Fe(s) + 4H₂O(g) ⇌ Fe₃O₄(s) + 4H₂(g) Note: The reaction with iron is reversible, and the product is iron(II, III) oxide, also known as magnetite.

{{VISUAL: diagram: laboratory apparatus to show the reaction of a metal with steam. It includes a flask of boiling water to generate steam, which passes over a heated metal sample (like iron wool) in a horizontal test tube, with the resulting hydrogen gas collected over water in an inverted test tube.}}

{{KEY: exam | title=Hydroxide vs. Oxide | text=A very common question asks for the product of a metal's reaction with water. Remember: Highly reactive metals (Na, K, Ca) with liquid water form hydroxides. Less reactive metals (Al, Zn, Fe) with steam form oxides.}}

5. The Non-Reactors

Finally, we have the metals at the bottom of the reactivity series. Metals like Lead (Pb), Copper (Cu), Silver (Ag), and Gold (Au) do not react with water or even steam. This is why they are often used for making water pipes (historically, lead), hot water tanks (copper), and jewellery (silver, gold), as they resist corrosion by water.

The intensity of a metal's reaction with water—from explosive reactions with cold water to no reaction at all—perfectly mirrors its position in the reactivity series. The higher the metal, the more vigorous the reaction.

What happens when Metals react with Acids?

{{FORMULA: expr=Metal + Dilute Acid → Metal Salt + Hydrogen Gas | symbols=→:yields}}

What Happens When Metals React with Acids?

In our journey through the properties of metals, we've seen how they react with air and water. Now, let's explore one of their most characteristic chemical properties: their reaction with acids. Have you ever wondered why we are advised not to store sour substances like curd, pickles, or lemon juice in metal containers (like those made of copper or brass)? The answer lies in this very reaction!

Most metals react with dilute acids to produce a metal salt and hydrogen gas. This is a classic example of a displacement reaction, where the more reactive metal displaces the less reactive hydrogen from the acid.

The General Reaction and Formation of Hydrogen Gas

The general word equation for this reaction is simple and powerful:

Metal + Dilute Acid → Metal Salt + Hydrogen Gas

Let's take the example of zinc metal reacting with dilute hydrochloric acid (HCl). Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)

Here, zinc (Zn) displaces hydrogen from HCl to form zinc chloride (ZnCl₂), which is a salt, and liberates hydrogen gas (H₂). The same happens with dilute sulfuric acid (H₂SO₄):

Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)

This reaction is often demonstrated in the lab. It's quite lively, with bubbles of hydrogen gas fizzing out from the surface of the metal.

{{VISUAL: photo: laboratory setup showing zinc granules in a test tube with dilute sulfuric acid being added. A delivery tube channels the gas produced through a trough of soap solution, creating bubbles filled with hydrogen gas.}}

So, we see bubbles, but how can we be sure it's hydrogen gas?

The "Pop" Test: Confirming Hydrogen Gas

There's a definitive and famous test for hydrogen gas. Because hydrogen is highly flammable, it burns with a characteristic sound when exposed to a flame.

Procedure:

1. Collect the gas being evolved from the reaction in a test tube.
2. Bring a burning candle or a matchstick near the mouth of the test tube.
3. If hydrogen gas is present, you will hear a distinct 'pop' sound as the gas burns rapidly.

This simple test is a reliable confirmation of the presence of hydrogen gas.

{{KEY: points | title=Testing for Hydrogen Gas | text=- The reaction of a metal with a dilute acid produces a gas.

• This gas is passed through a soap solution to form bubbles for easy handling.
• When a burning splint is brought near a gas-filled bubble, it burns with a characteristic 'pop' sound.
• This 'pop' sound confirms the presence of hydrogen gas.}}

Reactivity Series: The Deciding Factor

Does every metal react with acids in the same way? Not at all! The vigour and speed of the reaction depend entirely on the metal's position in the reactivity series.

• Highly Reactive Metals: Metals like Potassium (K), Sodium (Na), and Calcium (Ca) are at the top of the series. They react explosively with dilute acids. These reactions are too violent to be performed safely in a school lab.

• Moderately Reactive Metals: Metals like Magnesium (Mg), Aluminium (Al), Zinc (Zn), and Iron (Fe) react readily with dilute acids. The rate of bubble formation (effervescence) shows their relative reactivity. Magnesium reacts most vigorously, followed by aluminium, then zinc, and finally iron, which reacts quite slowly.

• Less Reactive Metals: Metals that are below hydrogen in the reactivity series, such as Copper (Cu), Mercury (Hg), Silver (Ag), and Gold (Au), cannot displace hydrogen from dilute acids. Therefore, they do not react with them.

Cu(s) + HCl(aq) → No reaction

This is precisely why copper vessels are used for many purposes, but not for storing acidic foods for long periods, as other slow reactions can still occur.

{{VISUAL: chart: a simplified reactivity series from Potassium to Gold. Arrows on the side show 'Decreasing Reactivity'. Icons next to Mg, Zn, Fe, and Cu show their reaction with dilute HCl: 'Vigorous Bubbling', 'Steady Bubbling', 'Slow Bubbling', and a 'No Reaction' cross mark respectively.}}

{{KEY: concept | title=The Role of the Reactivity Series | text=The reactivity series is a list of metals arranged in order of their decreasing chemical reactivity. A metal placed higher in the series can displace a metal (or hydrogen) that is lower in the series from its salt solution (or acid). This principle governs whether a reaction between a metal and an acid will occur and how vigorous it will be.}}

An Important Exception: The Case of Nitric Acid

There's a fascinating twist when metals react with nitric acid (HNO₃). You might expect hydrogen gas, but in most cases, it isn't produced. Why?

Nitric acid is a very strong oxidizing agent. As soon as hydrogen gas (H₂) is formed in the reaction, the HNO₃ immediately oxidizes it to form water (H₂O). Meanwhile, the nitric acid itself gets reduced to one of the nitrogen oxides, such as N₂O, NO, or NO₂, depending on the concentration of the acid and the nature of the metal.

3Cu + 8HNO₃ (dilute) → 3Cu(NO₃)₂ + 2NO + 4H₂O Notice: Water is formed, not hydrogen gas!

{{ZOOM: title=Aqua Regia: The "Royal Water" | text=Even highly unreactive metals like gold (Au) and platinum (Pt) can be dissolved. They don't react with individual acids, but they do react with a freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in a 3:1 ratio. This mixture is called Aqua Regia (Latin for "royal water") because of its ability to dissolve noble metals.}}

However, there are two exceptions to the nitric acid rule! Magnesium (Mg) and Manganese (Mn) are the only two metals that react with very dilute nitric acid to produce hydrogen gas.

Mg(s) + 2HNO₃(very dilute) → Mg(NO₃)₂(aq) + H₂(g)

{{KEY: exam | title=The Nitric Acid Anomaly | text=Questions based on the reaction of metals with nitric acid are a favourite in exams. Always remember to state that HNO₃ is a strong oxidizing agent that oxidizes the H₂ produced to H₂O. Do not forget to mention the two exceptions: Magnesium (Mg) and Manganese (Mn) with very dilute HNO₃.}}

Key Takeaway: The reaction between a metal and an acid is a powerful demonstration of the metal's reactivity, governed by its place in the reactivity series, with the unique behaviour of nitric acid serving as a crucial exception.