Metals and Non-metals

Metals and Non-metals

Metals

1. Metals are a group of elements known for several characteristic properties. They are good conductors of both heat and electricity, and they possess malleability and ductility. Metals typically have a shiny appearance (lustrous) and are hard, strong, heavy, and produce a ringing sound when struck (sonorous).
2. Some common examples of metals include iron, aluminium, copper, silver, gold, platinum, zinc, tin, lead, sodium, potassium, calcium, and magnesium. With the exception of mercury, which is a liquid at room temperature, all metals are solid.
3. During chemical reactions, metals have the ability to lose electrons and form positively charged ions, making them electropositive elements.
4. Metals find widespread use in various applications, including construction, transportation, electronics, energy generation, manufacturing, and medicine. They are vital for national economies and are abundant in the Earth's crust, with elements like aluminium and iron being among the most prevalent.

Non-metals

1. Non-metals are elements that exhibit properties contrasting to those of metals. They are generally poor conductors of heat and electricity and lack malleability and ductility. Non-metals are not shiny (non-lustrous) and have a dull appearance. They are often soft, weak and lightweight, and do not produce a ringing sound when struck (non-sonorous).
2. Examples of non-metals include carbon, sulphur, phosphorus, silicon, hydrogen, oxygen, nitrogen, chlorine, bromine, iodine, helium, neon, and argon.
3. Non-metals can form negative ions by gaining electrons during chemical reactions, earning them the label of electronegative elements. An exception is hydrogen, the only non-metal that can lose electrons to form positive ions (hydrogen ions or H⁺).
4. Non-metals are crucial for life on Earth, with carbon serving as the foundation of organic compounds necessary for living organisms. Oxygen is essential for respiration and combustion, while nitrogen's inertness makes air safer. Sulphur is present in various biological substances, and non-metals are used in the production of products like vegetable ghee, fertilisers, acids, explosives, and fungicides.
5. In terms of abundance in the Earth's crust, oxygen is the most abundant non-metal, followed by silicon, while elements like silicon, phosphorus, and sulphur are also significant components. Despite their smaller number compared to metals, non-metals play vital roles in the composition of Earth, its atmosphere, and oceans, making them essential for our planet's ecosystems and functioning.

Physical Properties of Metals and Non-metals

Metals

1. Malleability: Metals can be beaten into thin sheets without breaking. This property is known as malleability. For example, if you take a piece of aluminium metal and hammer it, it can be transformed into a thin aluminium sheet.
2. Ductility: Metals can be drawn or stretched into thin wires without breaking. This property is referred to as ductility. Common metals like copper, aluminium, and iron can be drawn into thin wires that are used in various applications, including electrical wiring.
3. Good Conductors of Heat: Metals are excellent conductors of heat. They allow heat to pass through them easily. Silver is the best conductor of heat among metals, followed by copper and aluminium. This property is essential in applications like cooking utensils and heat exchangers.
4. Good Conductors of Electricity: Metals are also good conductors of electricity. They allow electric current to flow through them with minimal resistance. Copper and aluminium are widely used in electrical wiring due to their excellent conductivity. Silver is the best conductor of electricity, but it is less common due to its cost.
5. Lustrous Appearance: Most metals have a shiny or lustrous surface when freshly exposed. This property, known as metallic lustre, makes metals visually appealing and suitable for decorative purposes. Gold, silver, and copper are renowned for their shiny appearance.
6. Hardness: Metals generally exhibit hardness and are resistant to deformation. This property ensures that metals maintain their structural integrity under various conditions. Iron, steel, and titanium are examples of hard metals used in construction and engineering.
7. Strength: Metals possess a high degree of mechanical strength, enabling them to withstand heavy loads without breaking or deforming. This strength makes metals suitable for structural components in buildings, bridges, and machinery.
8. State at Room Temperature: Most metals are solids at room temperature. Mercury is an exception, as it is the only metal that is liquid at room temperature.
9. High Melting and Boiling Points: Metals typically have high melting and boiling points. For instance, iron melts at around 1535^oC, and copper at about 1083^oC. These high melting points make metals suitable for high-temperature applications such as metalworking and smelting.
10. High Density: Metals have high densities, meaning they are heavy for their volume. For example, lead is a dense metal and is used for shielding against radiation due to its density.
11. Sonorousness: Metals produce a ringing sound when struck with an object. This property, called sonorousness, is used in making musical instruments like bells and cymbals.
12. Colour: Most metals have a silver or grey colour when freshly exposed. However, some metals, like copper, have distinct colours. Copper has a reddish-brown colour, while gold is characterised by its yellow hue.

Non-metals

1. Brittleness: Non-metals are brittle, which means they break into pieces when hammered, stretched, or subjected to mechanical stress. Unlike metals, they cannot be easily flattened into thin sheets or drawn into wires. Examples of brittle non-metals include sulphur, phosphorus, and carbon (in some forms).
2. Poor Conductors of Heat and Electricity: Non-metals do not conduct heat and electricity as effectively as metals. This is because they lack free electrons, which are necessary for electrical conductivity. While carbon in the form of graphite is an exception and is a good conductor of electricity, most non-metals, such as sulphur and phosphorus, are poor conductors of heat and electricity.
3. Dull Appearance: Non-metals generally have a dull or non-lustrous appearance. Unlike metals, they do not have a shiny surface. For example, sulphur and phosphorus are non-lustrous and appear dull. Iodine is an exception among non-metals, as it has a shiny surface similar to metals.
4. Softness: Most non-metals are relatively soft materials. They lack the hardness and mechanical strength seen in metals. An exception to this is diamond, which is an allotropic form of carbon and is extremely hard, making it one of the hardest known substances.
5. Physical State at Room Temperature: Non-metals can exist in various physical states at room temperature. Some are solid (e.g., sulphur, phosphorus, and carbon), while others are liquids (e.g., bromine), and some are gases (e.g., hydrogen, oxygen, nitrogen, and chlorine). Mercury is an exception among metals as it is a liquid at room temperature.
6. Low Melting and Boiling Points (Generally): Non-metals usually have lower melting and boiling points compared to metals. For instance, sulphur has a relatively low melting point of 115^oC. However, exceptions to this rule include carbon in the form of diamond, which has an extremely high melting and boiling point.
7. Low Density: Non-metals tend to have lower densities, meaning they are lighter for their volume compared to metals. For example, sulphur has a low density of 2 g/cm³.
8. Non-Sonorous: Non-metals do not produce a ringing sound when struck with an object. They lack the sonorous quality seen in metals.
9. Varied Colours: Non-metals exhibit a wide range of colours. For instance, sulphur is yellow, phosphorus can be white or red, and chlorine appears yellowish-green. Some non-metals, like hydrogen and oxygen, are colourless.

Exceptions in Physical Properties of Metals and Non-metals

The physical properties of metals and non-metals, as we've discussed, come with a variety of exceptions that make classification based solely on these properties challenging. Let's summarise these exceptions:

1. Electrical Conductivity: While non-metals are generally poor conductors of electricity, carbon in the form of graphite is an exception, displaying excellent electrical conductivity similar to metals.
2. Lustre: Typically, non-metals lack a shiny appearance (lustre). However, iodine stands out as an exception among non-metals, having a lustrous, reflective surface like metals.
3. Hardness and Softness: Although metals are usually hard, alkali metals like lithium, sodium, and potassium defy this trend and are soft, resembling non-metals in this aspect. Conversely, solid non-metals are commonly soft, but diamond, a non-metal, is an extraordinary exception with extreme hardness comparable to metals.
4. Physical State: Metals are typically found in the solid state at room temperature, but mercury, a metal, is a notable exception as it is a liquid at room temperature.
   Melting Points and Boiling Points: In general, metals have high melting and boiling points. However, sodium, potassium, gallium, and caesium are exceptions among metals, exhibiting low melting points, akin to non-metals. On the contrary, diamond, a non-metal, displays extremely high melting and boiling points, similar to metals.
5. Density: Metals are known for their high density, but alkali metals such as lithium, sodium, and potassium deviate from this pattern, having low densities resembling those of non-metals.

Chemical Properties of Metals and Non-metals

Metals

Chemical properties of metals are the characteristics that describe how metals react with other substances and elements. Metals display a wide range of chemical properties, which are a result of their tendency to lose electrons and form positively charged ions (cations). Here are some key chemical properties of metals:

1. Reactivity

Metals vary in their reactivity. Some metals, like sodium and potassium, are highly reactive and readily react with air and water, while others, like gold and platinum, are much less reactive and remain unreactive in most conditions.

Reactivity Series (or Activity Series) of Metals

1. The reactivity series is an arrangement of metals in a vertical column based on their chemical reactivity.
2. The most reactive metal is placed at the top of the series, and the least reactive metal is at the bottom.

Reactivity Trends

1. Metals at the top of the reactivity series are highly reactive, while those at the bottom are less reactive.
2. Highly reactive metals, like potassium and sodium, can react vigorously with substances like water and acids.
3. Less reactive metals, like copper and gold, do not readily react with water and acids.

Metals More Reactive Than Hydrogen

1. Metals placed above hydrogen in the reactivity series lose electrons more readily than hydrogen.
2. As a result, these metals are more reactive than hydrogen and can displace hydrogen from its compounds, such as water and acids, to produce hydrogen gas.
3. Examples of such metals include potassium, sodium, calcium, magnesium, and others in the list above.

Metals Less Reactive Than Hydrogen

1. Metals placed below hydrogen in the reactivity series lose electrons less readily than hydrogen.
2. These metals are less reactive than hydrogen and cannot displace hydrogen from its compounds like water and acids to form hydrogen gas.
3. Examples of less reactive metals are copper, mercury, silver, and gold.

2. Corrosion

Many metals corrode when exposed to oxygen and moisture. For example, iron corrodes to form iron oxide (rust).

Example: 4Fe (s) + 3O₂ (g) + 6H₂O (l) → 4Fe(OH)₃ (s) (Rust formation)

3. Alloy Formation

Alloys are mixtures of metals or metals with non-metals. They often have superior properties compared to pure metals. Brass, an alloy of copper and zinc, is stronger than pure copper.

Example: Brass - Cu (copper) + Zn (zinc)

4. Electrochemical Reactions

Metals are crucial in electrochemical processes, like batteries. In a zinc-carbon battery, zinc undergoes oxidation and provides electrical energy.

Example: Zn (s) + 2MnO₂ (s) + 2NH₄Cl (aq) → ZnCl₂ (aq) + 2MnO(OH) (s) + 2NH₃ (aq)

5. Catalytic Properties

Some metals, like platinum in catalytic converters, accelerate chemical reactions without being consumed themselves. In this case, platinum helps convert harmful exhaust gases into less harmful substances.

Example: Pt (platinum) catalysing the conversion of carbon monoxide (CO) to carbon dioxide (CO₂) in an automobile exhaust system.

6. Reaction of Metals with Oxygen

Most metals react with oxygen from the air to form metal oxides. This process is called oxidation.
The general equation for this reaction is:
Metal + Oxygen ⟶ Metal Oxide

Example: When iron reacts with oxygen, it forms iron oxide (rust): 4Fe + 3O₂ ⟶ 2Fe₂O₃ (Iron oxide or rust)

7. Reaction of Metals with Water

The reactivity of metals with water varies. Some metals react vigorously with cold water, while others react only with steam or not at all.
The general equation for the reaction of metals with water is:
Metal + Water ⟶ Metal Hydroxide + Hydrogen Gas

Example: Sodium reacts violently with cold water, producing sodium hydroxide and hydrogen gas: 2Na + 2H₂O ⟶ 2NaOH + H₂

8. Reaction of Metals with Dilute Acids

Most metals react with dilute acids, such as hydrochloric acid (HCl) or sulfuric acid (H₂SO₄), to produce salt and hydrogen gas.
The general equation for this reaction is: Metal + Dilute Acid → Salt + Hydrogen gas

Example: When zinc reacts with hydrochloric acid: Zn (s) + 2HCl (aq) → ZnCl₂ (aq) + H₂ (g)

Zinc displaces hydrogen from hydrochloric acid, forming zinc chloride and liberating hydrogen gas. This reaction is vigorous.
However, some metals like copper, silver, and gold do not react with dilute acids and remain unreactive.

9. Reaction of Metals with Salt Solutions

When a more reactive metal is placed in the salt solution of a less reactive metal, a displacement reaction occurs.
The more reactive metal displaces the less reactive metal from the salt solution, forming its own salt solution and releasing the less reactive metal.

Example: When zinc is placed in copper sulphate solution: CuSO₄ (aq) + Zn (s) → ZnSO₄ (aq) + Cu (s)

Zinc displaces copper from copper sulphate solution, forming zinc sulphate and depositing copper metal.

10. Reaction of Metals with Chlorine

Many metals react with chlorine gas (Cl₂) to form metal chlorides.
The general equation for this reaction is: Metal + Chlorine → Metal Chloride

Example: When iron reacts with chlorine: 2Fe (s) + 3Cl₂ (g) → 2FeCl₃ (s)

Iron combines with chlorine to form iron(III) chloride, a reddish-brown solid.

11. Reaction of Metals with Hydrogen

Some very reactive metals, like sodium, potassium, calcium, and magnesium, can react with hydrogen gas to form metal hydrides.

The general equation for this reaction is: Metal + Hydrogen → Metal Hydride

Example: When sodium reacts with hydrogen: 2Na (s) + H₂ (g) → 2NaH (s)

Sodium reacts with hydrogen to form sodium hydride, a solid compound.

Aqua-regia

Aqua-regia, often referred to as "royal water," is a highly corrosive and potent mixture of concentrated nitric acid (HNO₃) and concentrated hydrochloric acid (HCl). It is called "royal water" because of its ability to dissolve noble metals like gold and platinum, which are typically resistant to most acids. Here are some key characteristics and uses of aqua-regia:

1. Composition: Aqua-regia is prepared by mixing one part of concentrated nitric acid (HNO₃) with three parts of concentrated hydrochloric acid (HCl), typically in a laboratory setting. The resulting mixture has a distinctive reddish-brown colour.
2. Corrosive Nature: Aqua-regia is extremely corrosive and can cause severe chemical burns on contact with skin. It also releases toxic fumes, so it must be handled with great care in a well-ventilated area, preferably under a fume hood.
3. Versatility: Aqua-regia is known for its remarkable ability to dissolve a wide range of metals, including noble metals like gold and platinum. This makes it valuable in various chemical and metallurgical applications.

Non-metal

Chemical properties and reactions of non-metals are quite distinct from those of metals. Non-metals, being on the right side of the periodic table, generally have a tendency to gain or share electrons in chemical reactions. Here are the key chemical properties and reactions of non-metals:

1. Electronegativity

Non-metals generally have high electronegativity, meaning they have a strong tendency to attract electrons when they participate in chemical reactions. This property allows them to form covalent bonds by sharing electrons with other non-metals or with hydrogen.

2. Formation of Covalent Compounds

Non-metals often form covalent compounds, where atoms share electrons to achieve a stable electron configuration. In covalent bonds, electrons are shared rather than transferred as in ionic bonds.

3. Reaction with Oxygen

Non-metals often react with oxygen to form oxides. These oxides can be acidic or neutral.

Acidic oxides: Non-metals like carbon, sulphur, and nitrogen form acidic oxides (e.g., CO₂, SO₂). When these oxides dissolve in water, they produce acids (e.g., carbonic acid, sulphurous acid).
For example: C + O₂ → CO₂
CO₂ + H₂O → H₂CO₃ (Carbonic Acid)

Neutral oxides: Some non-metals like hydrogen and oxygen form neutral oxides (e.g., H₂O). These oxides do not affect the pH of water and are not acidic.

4. Reaction with Water

Non-metals do not generally react with water or steam to produce hydrogen gas (H₂). Unlike metals, they lack the ability to donate electrons to reduce water molecules into hydrogen gas and hydroxide ions (OH^-).

5. Reaction with Dilute Acids

Non-metals do not readily react with dilute acids. They do not displace hydrogen ions (H⁺) from acids because they are not good electron donors.

6. Reaction with Salt Solutions

In some cases, a more reactive non-metal can displace a less reactive non-metal from its salt solution. This displacement reaction forms new compounds.
For instance, chlorine gas (Cl₂) can displace bromine (Br₂) from a solution of sodium bromide (NaBr), resulting in the formation of sodium chloride (NaCl) and bromine gas (Br₂).
Cl₂ + NaBr → NaCl + Br₂

7. Reaction with Chlorine

Non-metals can react with chlorine gas (Cl₂) to form covalent chlorides. These chlorides are typically covalent compounds and do not conduct electricity.
Examples include hydrogen chloride (HCl) formed by the reaction of hydrogen (H₂) with chlorine (Cl₂) and phosphorus trichloride (PCl₃) formed by the reaction of phosphorus (P₄) with chlorine.
P₄ + 6Cl₂ → 4PCl₃

8. Reaction with Hydrogen

Non-metals can react with hydrogen gas (H₂) to form covalent hydrides. These hydrides involve the sharing of electrons and are typically covalent compounds.
Examples include hydrogen sulphide (H₂S) formed by sulphur's reaction with hydrogen and ammonia (NH₃) formed by nitrogen's reaction with hydrogen.
Non-metal hydrides do not contain ions and do not conduct electricity.

9. Reaction with Halogens

Non-metals can react with halogens like fluorine (F₂), chlorine (Cl₂), bromine (Br₂), and iodine (I₂) to form various covalent compounds.
For example, iodine reacts with hydrogen to form hydrogen iodide (HI).
Metals and non-metals react with each other differently due to their distinct chemical properties. The way they react depends on the number of electrons in their outermost energy levels (valence electrons) and their tendency to gain or lose electrons.

Cause of Chemical Bonding

A chemical bond is the force of attraction that holds atoms (or ions) together in a molecule or compound. It is the result of interactions between electrons in the outermost energy levels (valence electrons) of atoms.

Atoms tend to combine with each other because they seek to achieve a stable electron configuration, specifically, the electron arrangement of an inert gas (also known as a noble gas). Inert gases have a full outermost electron shell, which makes them highly stable and unreactive.

1. Methods to Achieve Stability: Atoms can achieve this stable electron configuration in three main ways:
2. Losing Electrons (Cations): Atoms can lose one or more electrons to other atoms, resulting in the formation of positively charged ions (cations). This happens when atoms have fewer than 8 electrons in their outermost shell (or fewer than 2 in the case of the first electron shell, which can hold a maximum of 2 electrons).
3. Gaining Electrons (Anions): Atoms can gain one or more electrons from other atoms, leading to the creation of negatively charged ions (anions). This occurs when atoms have more than 8 electrons (except for hydrogen and helium, which can be stable with just 2 electrons) in their outermost shell.
4. Sharing Electrons (Covalent Bonds): Atoms can share one or more electrons with other atoms. This is common among non-metals and results in the formation of covalent compounds. Sharing electrons allows atoms to complete their outermost electron shells and achieve stability.

How Do Metals and Non-metals React?

1. Ionic Compounds: When metals react with non-metals, they tend to form ionic compounds. Ionic compounds are held together by ionic bonds. These bonds are formed through the transfer of electrons from the metal atom (which loses electrons to become a positively charged ion or cation) to the non-metal atom (which gains electrons to become a negatively charged ion or anion). The electrostatic attraction between these oppositely charged ions holds them together in a stable lattice structure.
2. Covalent Compounds: When non-metals react with other non-metals, they typically form covalent compounds. Covalent compounds are held together by covalent bonds. In covalent bonds, atoms share electrons to achieve a stable electron configuration. This sharing of electrons allows both atoms to complete their outermost electron shells, resulting in a molecule.
3. Metals and Metals: Metals do not typically react with other metals to form compounds. This is because metals generally have low electronegativity and lose electrons easily to form cations. When two metals come into contact, they do not have a strong tendency to share or gain electrons from each other. Instead, metals often form alloys, which are mixtures of different metal elements.

Ions and Ionic Bond

Ions

Ions are electrically charged particles that result from the gain or loss of one or more electrons by an atom. Ions can be either positively charged (cations) or negatively charged (anions), depending on whether they have lost or gained electrons.

Formation of Ions

Cations (Positive Ions)

1. Cations are formed when an atom loses one or more electrons. This loss of electrons reduces the negative charge in the atom, leading to a net positive charge. Commonly, metals tend to lose electrons and form cations.
2. For example:
   Sodium (Na) has one valence electron. When it loses this electron, it becomes a sodium ion with a 1⁺ charge: Na⁺.
   Aluminium (Al) has three valence electrons. When it loses three electrons, it becomes an aluminium ion with a 3⁺ charge: Al³⁺.

Anions (Negative Ions)

1. Anions are formed when an atom gains one or more electrons. This gain of electrons increases the negative charge in the atom, resulting in a net negative charge. Non-metals typically gain electrons to form anions.
2. For example:
   Chlorine (Cl) has seven valence electrons. When it gains one electron, it becomes a chloride ion with a 1- charge: Cl^-.
   Oxygen (O) has six valence electrons. When it gains two electrons, it becomes an oxide ion with a 2- charge: O^2-.

Ionic Bonds

Ionic bonds are formed between ions, typically a cation and an anion. These bonds result from the electrostatic attraction between oppositely charged ions. Ionic bonds are characterised by the transfer of electrons from one atom to another, leading to the formation of a compound with a neutral overall charge. Here's how ionic bonds are formed:

Example 1: Sodium Chloride (NaCl)

Sodium (Na) is a metal and readily loses its one valence electron to form Na⁺.
Chlorine (Cl) is a non-metal and readily gains an electron to form Cl^-.

When sodium donates its electron to chlorine, they become oppositely charged ions (Na⁺ and Cl^-).
The electrostatic attraction between Na⁺ and Cl^- ions results in the formation of the ionic compound sodium chloride (NaCl), also known as table salt.

Example 2: Calcium Oxide (CaO)

Calcium (Ca) is a metal that loses two valence electrons to form Ca²⁺.
Oxygen (O) is a non-metal that gains two electrons to form O^2-.

When calcium and oxygen combine, Ca²⁺ and O^2- ions are attracted to each other, forming the ionic compound calcium oxide (CaO).

Ionic Compounds

Ionic compounds are chemical compounds formed through the combination of positively charged ions (cations) and negatively charged ions (anions). These compounds are also known as salts. Ionic compounds are typically composed of a metal cation and a non-metal anion, although there are exceptions.

Key Properties of Ionic Compounds:

1. High Melting and Boiling Points: Ionic compounds have high melting and boiling points because they are held together by strong electrostatic forces of attraction between oppositely charged ions. It takes a significant amount of energy to break these bonds and change the solid ionic lattice into a liquid or gas.
2. Solubility in Water: Many ionic compounds are soluble in water because water molecules can surround and separate the individual ions due to their polar nature. This allows ions to move freely in the aqueous solution, making the solution conductive.
3. Electrical Conductivity: Ionic compounds are poor conductors of electricity in their solid state but become good conductors when melted (molten) or dissolved in water. In these states, ions are free to move and carry electric charge, allowing the flow of electric current.
4. Crystalline Structure: Ionic compounds typically form regular, three-dimensional crystal lattices. The repeating pattern of ions in the crystal lattice contributes to the compound's stability and hardness.
5. Brittle Nature: Ionic compounds are often brittle, meaning they can break or shatter easily when subjected to mechanical stress. This brittleness is due to the arrangement of ions in the crystal lattice. When force is applied, like-charged ions align and repel each other, causing the crystal to fracture.
6. High Solubility in Polar Solvents: Apart from water, ionic compounds tend to be soluble in other polar solvents, such as ammonia and methanol, which can also break apart the crystal lattice by surrounding the ions.
7. Non-Volatile: Ionic compounds are generally non-volatile at room temperature, meaning they do not readily evaporate into a gas phase.
8. Colour: Many ionic compounds are white or colourless, but some may exhibit colour due to the presence of transition metal ions with partially filled d-orbitals. These transition metal ions can absorb specific wavelengths of light, giving rise to coloured compounds.
9. Hygroscopic Nature: Some ionic compounds are hygroscopic, meaning they readily absorb water vapour from the atmosphere and can even dissolve in the absorbed water, forming solutions.
10. Examples:
    Sodium Chloride (NaCl): Table salt.
    Calcium Carbonate (CaCO₃): Found in limestone and shells.
    Potassium Nitrate (KNO₃): Used in fertilisers and fireworks.
    Magnesium Sulphate (MgSO₄): Epsom salt.
    Ammonium Chloride (NH₄Cl): Used in dry cell batteries and as a flux in soldering.