Metals and Non-metals Class 10 Science Notes

Physical Properties of Metals

Metals are a class of elements that typically exhibit a distinct set of physical characteristics. These properties arise from the way their atoms are structured and bond together, particularly the presence of free electrons.

Key Physical Properties of Metals:

Property	Description	Example/Application
Lustre	Metals, in their pure state, possess a shiny surface. This characteristic appearance is known as metallic lustre. This shine is due to the reflection of light from the free electrons on the surface of the metal.	This property makes metals useful for decorative purposes. For instance, gold and silver are used in jewellery and ornaments because of their inherent shine.
Hardness	Most metals are hard and resist scratching or indentation. The degree of hardness varies significantly among different metals.	The hardness of metals makes them suitable for structural applications. Iron is used in construction, and steel (an alloy of iron) is used to make tools, machinery, and buildings. However, some metals like sodium and potassium are exceptions and are exceptionally soft.
Malleability	Malleability is the ability of a metal to be beaten or rolled into thin sheets without breaking. This property is a direct consequence of the ability of metal atoms to slide past each other under applied force, facilitated by the metallic bond.	Gold is the most malleable metal, meaning it can be hammered into extremely thin foils (gold leaf) used for gilding decorative objects. Aluminium is also highly malleable and is used to make aluminium foil for packaging food.
Ductility	Ductility refers to the ability of a metal to be drawn into thin wires. Similar to malleability, this property allows metal atoms to deform without fracturing.	Copper is widely used for electrical wiring due to its excellent ductility and conductivity. Gold is also exceptionally ductile; one gram of gold can be drawn into a wire about 2 kilometers long.
Conductivity	Metals are excellent conductors of heat and electricity. This is because the free, delocalized electrons in the metallic lattice can easily absorb thermal energy and move freely to carry electric charge.	Copper and aluminium are used in electrical wiring. Silver is the best conductor of heat and electricity, but its cost limits its widespread use. Metals like iron are used in cooking vessels because they conduct heat effectively to cook food. Lead and mercury are comparatively poor conductors.
Sonorousness	Metals are sonorous, meaning they produce a ringing sound when struck against a hard surface. This sound is produced by the vibration of the metal's lattice structure.	The sonorous property of metals is utilized in making school bells, gongs, and musical instruments like cymbals.
State at Room Temp.	Most metals exist as solids at room temperature. This is due to strong interatomic forces. The primary exception is mercury, which is a liquid at room temperature.	The solid state of most metals makes them suitable for construction and manufacturing. Mercury's liquid state at room temperature makes it useful in thermometers and barometers.
Melting & Boiling Points	Metals generally have high melting and boiling points. This indicates strong metallic bonds that require a significant amount of energy to break.	High melting points are crucial for metals used in high-temperature applications, like engine parts or furnace components. However, some metals like gallium and caesium have very low melting points and can melt in one's hand.

Physical Properties of Non-metals

Non-metals are elements that generally do not exhibit the characteristic properties of metals. They are found in various states at room temperature and have a different atomic structure.

Property	Description	Example/Application
Lustre	Non-metals are generally dull and do not have a shiny surface. They lack the free electrons that cause metallic lustre.	Most non-metals do not have a lustrous appearance. An exception is iodine, which is a non-metal but has a lustrous crystalline appearance.
Hardness	Non-metals vary greatly in hardness. Some are very soft (like sulphur), while others are very hard.	Diamond, an allotrope of carbon, is an exception as it is the hardest naturally occurring substance known. Carbon in other forms (like coal) is brittle.
Malleability	Non-metals are brittle and cannot be beaten into thin sheets. If hammered, they tend to break or shatter.	Non-metals are not used for applications requiring malleability.
Ductility	Non-metals are not ductile and cannot be drawn into wires.	Non-metals are not used for applications requiring ductility.
Conductivity	Non-metals are generally poor conductors of heat and electricity. They lack the free electrons that facilitate the movement of heat and charge.	Non-metals are often used as insulators. For example, the outer coating of electrical wires is made of non-conducting materials like PVC or rubber. An exception is graphite, an allotrope of carbon, which is a good conductor of electricity.
Sonorousness	Non-metals are not sonorous. They do not produce a ringing sound when struck.	Non-metals are not used for making bells or musical instruments requiring a resonant sound.
State at Room Temp.	Non-metals can exist as solids, liquids, or gases at room temperature.	Oxygen and hydrogen are gases. Bromine is a liquid. Carbon, sulphur, and phosphorus are solids.
Melting & Boiling Points	Non-metals generally have low melting and boiling points. Their interatomic forces are weaker compared to metals.	Many non-metals vaporize easily. However, diamond (an allotrope of carbon) has a very high melting and boiling point due to its strong covalent network structure.

Remember
While these are general properties, there are several exceptions like iodine (lustrous non-metal), graphite (conductive non-metal), and mercury (liquid metal) that are crucial to remember for exams.

Exam-smart
A question might present a list of elements and ask to identify metals or non-metals based on descriptions of their physical properties. For example, if a substance is described as being "drawn into wires and conducts electricity well," you would identify it as a metal. Conversely, a substance described as "brittle and dull" would likely be a non-metal. You might also be asked to explain why metals are used in specific applications (e.g., why copper is used for electrical wires).

Chemical Properties of Metals

Metals exhibit a range of chemical behaviours, primarily due to their tendency to lose electrons and form positive ions. This property influences their reactions with oxygen, water, acids, and solutions of other metal salts, leading to the formation of various compounds and observable changes. Understanding these reactions is crucial for predicting metal behaviour and for applications like corrosion prevention and metal extraction.

Reactions of Metals with Air

Metals react with oxygen present in the air, forming metal oxides. The reactivity of metals with oxygen varies significantly.

Vigorous Reaction: Highly reactive metals like potassium (K) and sodium (Na) react so violently with oxygen that they can catch fire. To prevent this, they are stored immersed in kerosene oil.
$2 K (s) + \frac{1}{2} O_{2} (g) \to K_{2} O (s)$
$2 N a (s) + \frac{1}{2} O_{2} (g) \to N a_{2} O (s)$
Formation of Oxide Layers: Metals such as magnesium (Mg), aluminium (Al), zinc (Zn), and lead (Pb) form a thin layer of their respective oxides on their surface at room temperature. This protective oxide layer prevents further reaction of the metal with oxygen.
$2 M g (s) + O_{2} (g) \to 2 M g O (s)$
$4 A l (s) + 3 O_{2} (g) \to 2 A l_{2} O_{3} (s)$
Burning: While iron filings burn vigorously when sprinkled into a flame, bulk iron does not burn on heating.
$3 F e (s) + 2 O_{2} (g) \overset{f l a m e}{\to} F e_{3} O_{4} (s)$
Surface Coating: Metals like copper (Cu) do not burn but get coated with a black layer of copper(II) oxide when heated strongly in air.
$2 C u (s) + O_{2} (g) \overset{h e a t}{\to} 2 C u O (s)$
No Reaction: Very unreactive metals like silver (Ag) and gold (Au) do not react with oxygen even at high temperatures.

Amphoteric Oxides: Some metal oxides, like aluminium oxide ( $A l_{2} O_{3}$ ) and zinc oxide ( $Z n O$ ), exhibit both acidic and basic properties. They are called amphoteric oxides because they react with both acids and bases to produce salt and water.

Reaction with acid: $A l_{2} O_{3} (s) + 6 H C l (a q) \to 2 A l C l_{3} (a q) + 3 H_{2} O (l)$
Reaction with base: $A l_{2} O_{3} (s) + 2 N a O H (a q) \to 2 N a A l O_{2} (a q) + H_{2} O (l)$ (Sodium aluminate)

Solubility of Metal Oxides: Most metal oxides are insoluble in water. However, oxides of highly reactive metals like sodium and potassium dissolve in water to form alkalis (soluble bases).

$N a_{2} O (s) + H_{2} O (l) \to 2 N a O H (a q)$
$K_{2} O (s) + H_{2} O (l) \to 2 K O H (a q)$

Anodising: This is a process used to form a thick, protective oxide layer on the surface of metals like aluminium. The aluminium article is made the anode in an electrolytic cell with dilute sulphuric acid. Oxygen evolved at the anode reacts with aluminium to form a thicker aluminium oxide ( $A l_{2} O_{3}$ ) layer, which increases resistance to corrosion and can be dyed for aesthetic appeal.

Remember
Sodium and potassium are stored under kerosene to prevent their reaction with oxygen and moisture in the air.

Exam-smart
A question might provide an observation like "a metal burns with a dazzling white flame," and you'll need to identify the metal (e.g., Magnesium) or the product formed. It could also ask about the properties of amphoteric oxides, requiring you to recall examples and their dual reactivity.

Reactions of Metals with Water

Metals react with water to produce a metal oxide and hydrogen gas ( $H_{2}$ ). If the metal oxide is soluble in water, it forms a metal hydroxide. The reactivity with water varies:

Violent Reaction with Cold Water: Highly reactive metals like potassium (K) and sodium (Na) react explosively with cold water, producing hydrogen gas and a lot of heat. The heat evolved is sufficient to ignite the hydrogen gas, causing fire.
$2 K (s) + 2 H_{2} O (l) \to 2 K O H (a q) + H_{2} (g) + heat$
$2 N a (s) + 2 H_{2} O (l) \to 2 N a O H (a q) + H_{2} (g) + heat$
Less Violent Reaction with Cold Water: Calcium (Ca) reacts less vigorously with cold water. The heat produced is not enough to ignite the hydrogen gas. Calcium starts to float because the hydrogen bubbles stick to its surface.
$C a (s) + 2 H_{2} O (l) \to C a (O H)_{2} (a q) + H_{2} (g)$
Reaction with Hot Water: Magnesium (Mg) does not react with cold water but reacts with hot water to form magnesium hydroxide and hydrogen gas. It also floats due to hydrogen bubbles.
$M g (s) + 2 H_{2} O (h o t) \to M g (O H)_{2} (a q) + H_{2} (g)$
Reaction with Steam: Metals like aluminium (Al), iron (Fe), and zinc (Zn) do not react with cold or hot water but react with steam to form metal oxides and hydrogen gas.
$2 A l (s) + 3 H_{2} O (g) \overset{s t e a m}{\to} A l_{2} O_{3} (s) + 3 H_{2} (g)$
$3 F e (s) + 4 H_{2} O (g) \overset{s t e a m}{\to} F e_{3} O_{4} (s) + 4 H_{2} (g)$
No Reaction with Water: Unreactive metals such as lead (Pb), copper (Cu), silver (Ag), and gold (Au) do not react with water or steam at all.

Remember
Metals that float when reacting with water are those that produce hydrogen gas, which adheres to the metal's surface, making it buoyant.

Exam-smart
Questions may ask to predict observations when a specific metal (e.g., Calcium) reacts with water, or to arrange metals based on their reactivity with different forms of water (cold water, hot water, steam). You might be given reactions and asked to identify the metal or the state of water used.

Reactions of Metals with Acids

Reactive metals react with dilute acids to produce a salt and hydrogen gas ( $H_{2}$ ). This reaction demonstrates the reducing nature of metals.

General Reaction:
Metal + Dilute Acid $\to$ Salt + Hydrogen gas
For example, with dilute hydrochloric acid ( $H C l$ ):
$M g (s) + 2 H C l (a q) \to M g C l_{2} (a q) + H_{2} (g)$
$A l (s) + 3 H C l (a q) \to A l C l_{3} (a q) + \frac{3}{2} H_{2} (g)$ (or $2 A l + 6 H C l \to 2 A l C l_{3} + 3 H_{2}$ )
$Z n (s) + 2 H C l (a q) \to Z n C l_{2} (a q) + H_{2} (g)$
$F e (s) + 2 H C l (a q) \to F e C l_{2} (a q) + H_{2} (g)$
Reactivity Order with Acids: The rate of hydrogen gas evolution indicates the reactivity of the metal. More vigorous bubbling means higher reactivity. The reactivity order observed with dilute $H C l$ is generally: $M g > A l > Z n > F e$ .
Magnesium shows the most vigorous reaction and the highest temperature increase, indicating it is the most reactive among these.
Copper (Cu) does not react with dilute hydrochloric acid, showing no bubbles and no temperature change. This implies copper is less reactive than hydrogen.
Reaction with Nitric Acid ( $H N O_{3}$ ): Nitric acid ( $H N O_{3}$ ) is a strong oxidizing agent. When metals react with dilute $H N O_{3}$ , hydrogen gas is usually not evolved because $H N O_{3}$ oxidizes the $H_{2}$ produced to water and itself gets reduced to various nitrogen oxides ( $N_{2} O$ , $N O$ , $N O_{2}$ ).
Exception: Metals like magnesium (Mg) and manganese (Mn) can react with very dilute nitric acid to evolve $H_{2}$ gas.
Aqua Regia: This is a freshly prepared, highly corrosive mixture of concentrated hydrochloric acid ( $H C l$ ) and concentrated nitric acid ( $H N O_{3}$ ) in a 3:1 ratio. Aqua regia is one of the few substances that can dissolve noble metals like gold (Au) and platinum (Pt), which do not react with individual acids.

Remember
The reaction of metals with acids produces hydrogen gas, unless the acid is a strong oxidizing agent like nitric acid, or the metal is unreactive (like copper).

Exam-smart
You might be given a scenario describing the effervescence (bubbling) observed when different metals are added to dilute acid, and asked to arrange the metals in order of reactivity or identify the most reactive one. Questions could also test your knowledge about why hydrogen is not evolved with nitric acid, or the composition and use of aqua regia.

How do Metals react with Solutions of other Metal Salts?

Reactive metals can displace less reactive metals from their compounds dissolved in water or in a molten state. This is a type of displacement reaction.

Principle: If metal A can displace metal B from its salt solution, then metal A is more reactive than metal B.
Metal A + Salt solution of B $\to$ Salt solution of A + Metal B
Example (Activity 3.12):
When a copper wire is placed in iron sulphate ( $F e S O_{4}$ ) solution: No reaction occurs because copper is less reactive than iron.
When an iron nail is placed in copper sulphate ( $C u S O_{4}$ ) solution: A reaction occurs. Iron displaces copper from copper sulphate solution, forming iron(II) sulphate ( $F e S O_{4}$ ) solution and depositing copper metal.
$F e (s) + C u S O_{4} (a q) \to F e S O_{4} (a q) + C u (s)$
This is a displacement reaction because iron displaces copper.
Application: This reaction type is fundamental to constructing the reactivity series of metals. By observing which metal displaces another from its salt solution, we can infer their relative reactivity.

Remember
A more reactive metal will displace a less reactive metal from its salt solution. If a metal does not displace another metal, it is less reactive than that metal.

Exam-smart
Case study questions often present a table showing observations of different metals (A, B, C, D) being added to various metal salt solutions. You will need to use these observations to determine the relative reactivity of the metals, identify the most and least reactive, and predict further reactions. For example, if metal B displaces metal A from its salt solution, B is more reactive than A.

The Reactivity Series

The reactivity series (or activity series) is an ordered list of metals arranged in decreasing order of their reactivity. This series is determined by observing their reactions with oxygen, water, acids, and their ability to displace less reactive metals from their salt solutions.

General Reactivity Series (Most Reactive to Least Reactive):

Order	Metal	Symbol	Reactivity
1	Potassium	K	Most
2	Sodium	Na
3	Calcium	Ca
4	Magnesium	Mg
5	Aluminium	Al
6	Zinc	Zn
7	Iron	Fe
8	Lead	Pb
9	Hydrogen	[H]	(Reference)
10	Copper	Cu	Least
11	Mercury	Hg
12	Silver	Ag
13	Gold	Au

Metals above Hydrogen: Metals like K, Na, Ca, Mg, Al, Zn, Fe, and Pb are more reactive than hydrogen. They react with dilute acids to liberate hydrogen gas.
Metals below Hydrogen: Metals like Cu, Hg, Ag, and Au are less reactive than hydrogen. They do not displace hydrogen from dilute acids.
Displacement Principle: Any metal can displace metals listed below it in the reactivity series from their salt solutions. Conversely, a metal cannot displace metals listed above it. For example, Iron (Fe) can displace Copper (Cu) from $C u S O_{4}$ solution because Fe is above Cu in the series, but Copper cannot displace Iron.

Remember
The reactivity series helps predict whether a displacement reaction will occur between a metal and a metal salt solution.

Exam-smart
You will likely be given a partial reactivity series or a set of displacement reactions and asked to arrange metals in order of reactivity, or predict the outcome of a reaction based on their position in the series. For instance, if you are told that metal X displaces metal Y from its salt solution, and metal Y displaces metal Z, you can deduce the order X > Y > Z.

How do metals and non-metals react?

Metals and non-metals react by transferring electrons to achieve a stable electron configuration, specifically a completely filled valence shell. This stability is similar to that of noble gases, which have very low chemical reactivity.

When a metal atom reacts with a non-metal atom, the metal atom tends to lose electrons from its outermost shell, forming a positively charged ion called a cation. The non-metal atom tends to gain these electrons into its outermost shell, forming a negatively charged ion called an anion.

For example, a sodium atom (Na) has one electron in its outermost (M) shell. By losing this electron, its L shell becomes the outermost shell with a stable octet (2, 8). This results in a sodium cation ( $N a^{+}$ ).

N a \to N a^{+} + e^{-}

(2, 8, 1) $\to$ (2, 8)

A chlorine atom (Cl) has seven electrons in its outermost shell. It needs one more electron to complete its octet. If it gains the electron lost by sodium, it forms a chloride anion ( $C l^{-}$ ).

C l + e^{-} \to C l^{-}

(2, 8, 7) $\to$ (2, 8, 8)

The oppositely charged ions, $N a^{+}$ and $C l^{-}$ , are attracted to each other by strong electrostatic forces of attraction. This attraction holds them together to form an ionic compound, such as sodium chloride (NaCl).

Formation of ionic compounds

Ionic compounds (also called electrovalent compounds) are formed by the transfer of electrons from a metal to a non-metal, resulting in the formation of oppositely charged ions that are held together by strong electrostatic forces.

Consider the formation of magnesium chloride ( $M g C l_{2}$ ): Magnesium (Mg) has two electrons in its outermost shell. It loses these two electrons to form a magnesium cation ( $M g^{2 +}$ ).

M g \to M g^{2 +} + 2 e^{-}

(2, 8, 2) $\to$ (2, 8)

Chlorine (Cl) needs one electron to complete its octet. To form $M g C l_{2}$ , two chlorine atoms are required, each accepting one electron.

C l + e^{-} \to C l^{-}

(2, 8, 7) $\to$ (2, 8, 8)

Therefore, one magnesium atom transfers two electrons, one to each of two chlorine atoms, forming $M g^{2 +}$ and two $C l^{-}$ ions. These ions then attract each other to form magnesium chloride.

Remember
Ionic compounds are formed between elements that readily lose electrons (metals) and elements that readily gain electrons (non-metals).

Exam-smart
Case studies may present electronic configurations of elements and ask to predict the type of compound formed and the ions present. For instance, if element X has configuration (2, 8, 1) and element Y has (2, 8, 7), they will form an ionic compound XY.

Properties of ionic compounds

Ionic compounds exhibit characteristic properties due to the strong electrostatic forces between their ions.

Property	Description	Reasoning
Physical nature	Ionic compounds are generally solids and are somewhat hard. They are also brittle and break into pieces when pressure is applied.	The strong electrostatic forces of attraction between the positive and negative ions hold them in a fixed lattice structure, giving them rigidity. However, when pressure is applied, layers of ions can shift, bringing like charges together, causing repulsion and fracture.
Melting and boiling points	Ionic compounds have high melting and boiling points. For example, NaCl melts at 1074 K and boils at 1686 K.	A large amount of energy is required to overcome the strong electrostatic forces of attraction between the ions in the crystal lattice to melt or boil the compound.
Solubility	Ionic compounds are generally soluble in water and insoluble in non-polar solvents like kerosene and petrol.	Water is a polar solvent. The polar water molecules can surround the ions, overcoming the attractive forces between them and hydrating the ions, thus dissolving the ionic compound. Non-polar solvents cannot effectively solvate the charged ions.
Conduction of electricity	Ionic compounds conduct electricity in the molten state or when dissolved in water, but not in the solid state.	In the solid state, ions are held in fixed positions in the lattice and cannot move. In the molten state or in aqueous solution, the ions are free to move and carry electric charge when an electric field is applied, thus conducting electricity.

Remember
The conduction of electricity by ionic compounds in molten or aqueous states is due to the mobility of ions.

Exam-smart
A question might describe an unknown compound's properties (e.g., dissolves in water, conducts electricity when molten) and ask the student to identify it as ionic or covalent, or to predict its behavior under different conditions.

Occurrence of Metals

The Earth's crust and seawater are the primary sources for obtaining metals. Elements or compounds that naturally occur in the Earth's crust are called minerals. When a mineral contains a very high percentage of a particular metal and that metal can be extracted profitably, it is then referred to as an ore.

The reactivity series of metals helps us understand how metals are extracted from their ores. Based on their reactivity, metals can be broadly categorized into three groups, influencing their occurrence and extraction methods:

Metals of low reactivity: These are the least reactive metals and are often found in the Earth's crust in their free state (as elements). Examples include gold, silver, and platinum. Copper and silver can also be found in combined states as sulphides or oxides.
Metals of medium reactivity: These metals are moderately reactive. They are found mainly as oxides, sulphides, or carbonates. Examples include zinc, iron, and lead. Their ores, if sulphides or carbonates, are usually converted to oxides before extraction.
Metals of high reactivity: These are very reactive metals and are never found in the Earth's crust as free elements. They are always found in their combined states, typically as oxides, chlorides, or sulphates. Examples include potassium, sodium, calcium, magnesium, and aluminium.

Extraction of Metals

Extraction of metals from their ores involves several steps, with the specific method depending on the metal's position in the reactivity series.

Enrichment of Ores

Ores mined from the Earth usually contain impurities like soil and sand, collectively known as the gangue. Removing these impurities is crucial and is called enrichment or dressing of the ore. The techniques used for enrichment exploit the differences in physical or chemical properties between the ore and the gangue.

Extracting Metals Low in the Activity Series

Metals at the bottom of the reactivity series are very unreactive. Their oxides can often be reduced to the pure metal by simply heating.

Example: Mercury Cinnabar (HgS) is an ore of mercury. When heated in air, it first forms mercuric oxide (HgO), which then decomposes into mercury and oxygen upon further heating.

2 HgS(s) + 3 O_{2} (g) \overset{Heat}{\to} 2 HgO(s) + 2 {SO}_{2} (g)

2 HgO(s) \overset{Heat}{\to} 2 Hg(l) + O_{2} (g)

Example: Copper Copper, found as $C u_{2} S$ , can also be extracted by heating.

2 {Cu}_{2} S(s) + 3 O_{2} (g) \overset{Heat}{\to} 2 {Cu}_{2} O(s) + 2 {SO}_{2} (g)

2 {Cu}_{2} O(s) + {Cu}_{2} S(s) \overset{Heat}{\to} 6 Cu(s) + {SO}_{2} (g)

Remember
For metals low in the reactivity series, direct heating of their oxides or sulphides can yield the metal.

Exam-smart
This method is applicable to obtaining metals from their oxides by reduction using heat. You might be given a reaction and asked to identify the ore and the metal obtained.

Extracting Metals in the Middle of the Activity Series

Metals like iron, zinc, lead, and copper are moderately reactive. They are usually found as sulphides or carbonates. Since it is easier to reduce metal oxides than sulphides or carbonates, these ores are first converted into oxides.

Roasting: Sulphide ores are strongly heated in the presence of excess air to convert them into metal oxides.

2 ZnS(s) + 3 O_{2} (g) \overset{Heat}{\to} 2 ZnO(s) + 2 {SO}_{2} (g)

Calcination: Carbonate ores are heated strongly in limited air (or absence of air) to convert them into metal oxides.

{ZnCO}_{3} (s) \overset{Heat}{\to} ZnO(s) + {CO}_{2} (g)

The metal oxides are then reduced to metals. Common reducing agents include carbon.

ZnO(s) + C(s) \to Zn(s) + CO(g)

Alternatively, more reactive metals can be used as displacing agents for metals of lower reactivity.

{3MnO}_{2} (s) + 4 Al(s) \to 3 Mn(l) + 2 {Al}_{2} O_{3} (s) + Heat

This reaction is highly exothermic, producing molten metal.

The reaction of iron(III) oxide ( $F e_{2} O_{3}$ ) with aluminium is known as the thermit reaction and is used to weld railway tracks or repair machinery:

F e_{2} O_{3} (s) + 2 Al(s) \to 2 Fe(l) + {Al}_{2} O_{3} (s) + Heat

Remember
Roasting is for sulphide ores, while calcination is for carbonate ores. Both aim to produce metal oxides.

Exam-smart
You might be given a chemical equation for roasting or calcination and asked to identify the process and the resulting product. Or, you could be given a thermit reaction and asked about its application.

Extracting Metals towards the Top of the Activity Series

Metals high up in the reactivity series (like sodium, potassium, calcium, magnesium, aluminium) are very reactive. Carbon is not a strong enough reducing agent to displace these metals from their oxides because they have a greater affinity for oxygen than carbon does.

These metals are obtained by electrolytic reduction of their molten salts or oxides.

Example: Sodium, Magnesium, Calcium They are obtained by the electrolysis of their molten chlorides. At the cathode (negative electrode): $N a^{+} + e^{-} \to N a$ At the anode (positive electrode): $2 C l^{-} \to C l_{2} + 2 e^{-}$
Example: Aluminium Aluminium is obtained by the electrolytic reduction of aluminium oxide ( $A l_{2} O_{3}$ ), usually dissolved in molten cryolite ( $N a_{3} A l F_{6}$ ) to lower the melting point.

Remember
Electrolysis of molten compounds is necessary for extracting highly reactive metals.

Exam-smart
A question might describe an electrolytic cell for a molten salt and ask you to identify the metal produced at the cathode or the gas produced at the anode.

Refining of Metals

Metals obtained through the reduction processes are often impure. Refining is the process of obtaining a pure metal from these impure metals. The most common method for refining is electrolytic refining.

Electrolytic Refining:

Setup: An impure metal is used as the anode (positive electrode) and a thin strip of the pure metal is used as the cathode (negative electrode). A solution of a salt of the metal is used as the electrolyte.
Process: When an electric current is passed, the pure metal from the anode dissolves into the electrolyte. An equivalent amount of pure metal from the electrolyte gets deposited on the cathode.
Impurities: Soluble impurities dissolve in the electrolyte, while insoluble impurities (like gold, silver, platinum) settle down at the bottom of the anode as anode mud.

This method is used for refining metals like copper, zinc, nickel, tin, silver, and gold.

Example: Electrolytic refining of Copper

Anode: Impure copper
Cathode: Strip of pure copper
Electrolyte: Acidified copper sulphate solution ( $C u S O_{4}$ solution with a little sulfuric acid)

When current flows, pure copper from the anode deposits on the cathode.

Remember
In electrolytic refining, the impure metal is the anode, the pure metal is the cathode, and the electrolyte is a salt solution of the metal. Anode mud contains precious metals.

Exam-smart
You could be shown a diagram of an electrolytic refining setup and asked to label the anode, cathode, and electrolyte, or to explain what happens to the impure metal and how pure metal is obtained.

Corrosion

Corrosion is the process where a metal is gradually destroyed by a chemical reaction with its environment. This usually involves oxidation.

Examples of Corrosion:

Silver: Tarnishes (turns black) when exposed to air due to the formation of a coating of silver sulphide ( $A g_{2} S$ ). This happens because silver reacts with sulphur ( $S$ ) present in the air.

4 A g + 2 H_{2} S + O_{2} \to 2 A g_{2} S + 2 H_{2} O

Copper: Loses its shiny brown surface and gains a green coat when it reacts with moist carbon dioxide ( $C O_{2}$ ) in the air. This green substance is basic copper carbonate ( $C u C O_{3} \cdot C u (O H)_{2}$ ). $2 C u + H_{2} O + C O_{2} + O_{2} \to C u C O_{3} \cdot C u (O H)_{2}$ (Green coating)
Iron: When exposed to moist air for a long time, it acquires a coating of a brown, flaky substance called rust. Rust is essentially hydrated iron(III) oxide.

Conditions for Rusting of Iron: Rusting of iron requires both moisture (water) and air (oxygen).

Activity 3.14 demonstrates this:
Test tube A: Iron nail in water exposed to air. Rusting occurs.
Test tube B: Iron nail in boiled distilled water with a layer of oil. Boiled water has less dissolved oxygen. Oil prevents further air dissolving. No rusting occurs.
Test tube C: Iron nail in dry air (using anhydrous calcium chloride to absorb moisture). No rusting occurs.

This shows that the presence of both oxygen and water is necessary for iron to rust.

Remember
Corrosion is an electrochemical process. Iron acts as the anode and gets oxidized to $F e^{2 +}$ ions, while oxygen acts as the cathode and gets reduced. Water facilitates the movement of ions.

Exam-smart
A case study might describe a scenario with iron objects in different environments (e.g., near the sea, in a dry warehouse, in a humid factory). You would need to apply the understanding of conditions required for rusting to predict where corrosion is most likely to occur.

Prevention of Corrosion

Corrosion can be prevented by stopping the metal from coming into contact with the corrosive environment. Common methods include:

Method	Description	How it Prevents Corrosion
Painting	Applying a coat of paint on the metal surface.	The paint layer acts as a barrier, preventing oxygen and moisture from reaching the metal surface.
Oiling/Greasing	Applying a layer of oil or grease on the metal surface.	Similar to painting, it creates a protective barrier against air and water. Commonly used for moving parts like machine components.
Galvanising	Coating the metal (usually iron or steel) with a thin layer of zinc ( $Z n$ ).	Zinc is more reactive than iron. It forms a protective layer of zinc oxide ( $Z n O$ ) and zinc carbonate ( $Z n C O_{3}$ ) which prevents further corrosion. Even if the zinc coating is scratched, the zinc preferentially corrodes (acts as a sacrificial metal) protecting the iron.
Chrome Plating	Coating the metal surface with a layer of chromium ( $C r$ ).	Chromium forms a tough, non-reactive layer of chromium oxide ( $C r_{2} O_{3}$ ) that protects the underlying metal.
Anodising	An electrolytic process used to increase the thickness of the natural oxide layer on the surface of metals like aluminium ( $A l$ ).	This forms a thicker, more protective oxide layer, making the metal more resistant to corrosion.
Alloying	Mixing the metal with other elements (metals or non-metals) to form an alloy.	Alloys can be designed to be much more resistant to corrosion than the pure metal. For example, stainless steel (iron with chromium and nickel) does not rust. Pure gold is soft, so it is alloyed with copper or silver to make jewellery harder and more durable.
Sacrificial Protection	Attaching a more reactive metal (e.g., magnesium ( $M g$ ), zinc ( $Z n$ )) to the metal to be protected (e.g., iron).	The more reactive metal corrodes preferentially, sacrificing itself to protect the less reactive metal. This is the principle behind galvanising and is also used for protecting underground pipelines.

What are Alloys? An alloy is a homogeneous mixture of two or more metals, or a metal and one or more non-metals. Alloys are prepared by melting the primary metal and dissolving other elements in it in definite proportions, followed by cooling.

Properties of Alloys:

Alloys generally have higher hardness and tensile strength than their constituent pure metals.
The electrical conductivity and melting point of an alloy are usually less than that of the pure metals.
Alloys can have unique properties that make them suitable for specific applications.

Examples of Alloys:

Brass: Alloy of copper ( $C u$ ) and zinc ( $Z n$ ). Used for making utensils, musical instruments.
Bronze: Alloy of copper ( $C u$ ) and tin ( $S n$ ). Used for making statues, medals, coins.
Stainless Steel: Alloy of iron ( $F e$ ), chromium ( $C r$ ), and nickel ( $N i$ ). Used for making utensils, surgical instruments.
Solder: Alloy of lead ( $P b$ ) and tin ( $S n$ ). Has a low melting point, used for welding electrical wires.

Amalgam: An alloy in which one of the metals is mercury ( $H g$ ). For example, sodium amalgam ( $N a amalgam$ ) is used in some chemical processes.

Remember
Stainless steel is resistant to rust because of the formation of a protective layer of chromium oxide on its surface.

Exam-smart
You might be given a situation where a metal object needs protection. You'll need to choose the most suitable prevention method based on the metal, its use, and cost-effectiveness, explaining why that method works (e.g., "Painting is suitable for an iron gate as it forms a barrier, but for underground pipes, sacrificial protection with magnesium is better because the paint layer would be damaged and zinc coating would not provide sacrificial protection in the absence of oxygen"). Questions about alloys will test your understanding of their composition and why they are preferred over pure metals.

Metals and Non-metals — Notes

Read the notes