CBSE Class 10 Science Chapter 3 Metals and Non-metals - Notes
Meaning of physical properties:
Physical properties are observable characteristics
of a substance that don't alter its chemical composition. They describe how a
material behaves physically, providing information about its structure and
behavior under different conditions.
Physical Properties of Metals and
Non-metals:
Physical Properties of Metals:
- Luster:
- Most metals have a shiny appearance when freshly polished. This property is often referred to as metallic luster. For Example: gold, silver and copper.
- Note: Some non-metals are lustrous, Like-Iodine and carbon. Carbon is lustrous only in certain forms like- diamond and graphite.
- Malleability:
- Metals can be hammered or rolled into thin sheets without breaking. This property is known as malleability.
- Ductility:
- Metals can be drawn into thin wires without fracturing. This property is called ductility.
- Conductivity:
- Metals are excellent conductors of heat and electricity. They allow the flow of heat and electric current through them with minimal resistance. For example: Copper, Graphite.
- High Density (except alkali metals):
- Metals are generally dense materials. They have a high mass per unit volume. Osmium (Os) has highest density which is 22.59 g/m3 and lithium (Li) has least density which is 0.534 g/m3.
- High Melting and Boiling Points (except Caesium (Cs) and Gallium (Ga)):
- Metals have high melting and boiling points. They can withstand high temperatures without changing their physical state.
- Note: Graphite which is a form of carbon, has a high boiling point and exists in a solid state at room temperature.
- Solid State (at Room Temperature):
- Most metals are solid at room temperature., with the exception of mercury, which is a liquid.
- Note: Some non-metal like- Bromine is a liquid at room temperature and Carbon and sulphur are solids at room temperature.
- Sonorous:
- Many metals produce a ringing or musical sound when struck, making them "sonorous".
- Hard and have a high tensile strength: metals are hard and have high tensile strength.
- Note: Carbon is the only non-metal with very high tensile strength.
- In Silver-grey in colour (except gold and copper):
- Metals usually find in silver or grey colour.
Physical Properties of Non-Metals:
- Lack of Luster:
- Non-metals often lack the characteristic metallic luster and appear dull or non-reflective.
- Brittle:
- Non-metals are generally brittle and can shatter or crumble when subjected to force or stress.
- Poor Conductors:
- Non-metals are poor conductors of heat and electricity. They do not allow the easy flow of heat and electric current.
- Low Density:
- Non-metals generally have lower density compared to metals. They have a lower mass per unit volume.
- Low Melting and Boiling Points:
- Non-metals typically have lower melting and boiling points compared to metals. They tend to be in the gaseous or liquid state at room temperature.
- Various States (at Room Temperature):
- Non-metals can exist in various states at room temperature, including solids like- sulfur and phosphorus, liquids like- bromine and gases like- oxygen and nitrogen.
- Non-Sonorous:
- Non-metals do not produce a sonorous sound when struck. They may produce a dull or non-musical sound.
- Lack of Malleability and Ductility:
- Non-metals are not malleable or ductile.They cannot be easily shaped into thin sheets or wires without breaking.
Exceptions in Physical Properties
- Alkali metals like- Na, K and Li are malleable enough to be cut with a knife.
- Mercury is a liquid metal at room temperature.
- Lead and mercury metals are surprisingly poor conductors of heat.
- Mercury show significant expansion even with slight temperature changes.
- Gallium and cesium metals are having exceptionally low melting points.
- Iodine which is a non-metal shows a lustrous appearance.
- Graphite which is non-metal conducts electricity.
- Diamond has extraordinarily high melting point and conduct heat efficiently.
Examples of Metals and Non-metals:
|
S.No. |
Metals |
Non-metals |
|
1 |
Iron
(Fe) – Solid |
Hydrogen
(H) – Gas |
|
2 |
Copper (Cu)– Solid |
Nitrogen (N) – Gas |
|
3 |
Aluminium
(Al)– Solid |
Oxygen
(O) – Gas |
|
4 |
Magnesium (Mg)– Solid |
Chlorine (Cl) – Gas |
|
5 |
Sodium
(Na)– Solid |
Fluorine
(F)– Gas |
|
6 |
Zinc (Zn)– Solid |
Bromine (Br)– Liquid |
|
7 |
Lead
(Pb)– Solid |
Iodine
(I) – Solid |
|
8 |
Gold (Au) – Solid |
Carbon (C) – Solid |
|
9 |
Silver
(Ag) – Solid |
Sulphur
(S) – Solid |
|
10 |
Tin (Sn) – Solid |
Phosphorous (P) – Solid |
|
11 |
Mercury
(Hg) – Liquid |
Silicon
(Si)– Solid |
Chemical Properties of Metals and Non-metals:
Chemical Properties of Metals:
Reaction of Metals with Oxygen (Metals
burnt in Air):
- When metals are burned and react with oxygen in the air formed Metal Oxides, which is classify as basic compounds find in nature.
- Metal Oxide can change the color of red litmus paper to blue.
Metal + Oxygen→ Metal oxide (basic)
Example: (i) when copper is heated in
air, it combines with oxygen and form copper (II) oxide (a black oxide).
2Cu + O2 → 2CuO
Example: (ii) When Aluminium react with
oxygen and form aluminium oxide.
4Al + 3O2 → 2Al2O3
Example: (iii) When K react with oxygen
and form Potassium oxide. It is a vigorous reaction.
4K(s)+O2(g)→2K2O(s)
Example: (iv) When Na react with oxygen
and form Sodium oxide. It is a vigorous reaction.
4Na(s)+O2(g)→2Na2O(s)
Note:
- Sodium and potassium are highly reactive metals.
- So, we must keep in them to kerosene oil to prevent reactions with oxygen, moisture and carbon dioxide in the air.
- The kerosene oil forms a protective layer, preventing direct contact with atmospheric gases and water, which could lead to explosions or fires.
2Mg(s)+O2(g)→2MgO(s)
4Al(s)+3O2(g)→2Al2O3(s)
Note: Silver, platinum and gold don’t
burn or react with air.
Basic Oxides of Metals:
- Basic oxides are also known as basic anhydrides.
- They are a type of chemical compound formed by metals and oxygen.
- These oxides have basic properties and can react with acids to form salts and water.
- They are normally found in crystalline solids state.
- Most metal oxide are insoluble in water but Some metallic oxides get dissolved in water and form alkalis and their aqueous solution turns red litmus blue.
- For Example: Sodium Oxide (Na2O), Potassium Oxide (K2O), Calcium Oxide (CaO), Magnesium Oxide (MgO), Barium Oxide (BaO), Strontium Oxide (SrO), Lithium Oxide (Li2O), Silver Oxide (Ag2O).
Na2O(s) + H2O(l) → 2NaOH(aq)
K2O(s) + H2O(l) → 2KOH(aq)
Amphoteric Oxides of Metals:
- Amphoteric oxides are metal oxides that can interact with both acids and bases, leading to the formation of salts and water.
- For Example: Aluminum Oxide (Al2O3), Zinc Oxide (ZnO), Lead(II) Oxide (PbO), Tin(II) Oxide (SnO), Chromium(III) Oxide (Cr2O3), Iron(III) Oxide (Fe2O3).
Al2O3(s) + 6HCl(aq) → 2AlCl3(aq)
+ 3H2O(l)
Al2O3(s) + 2NaOH(aq) → 2NaAlO2(aq)
+ H2O(l)
ZnO(s) + 2HCl(aq) → ZnCl2(aq) + H2O(l)
ZnO(s) + 2NaOH(aq) → Na2ZnO2(aq)
+ H2O(l)
Roasting:
- Roasting is a process of heating ores or metal sulfide compounds in the presence of air. This is done to drive off impurities and convert the ore into an oxide.
2ZnS(s) + 3O2(g) + Heat → 2ZnO(s) + 2SO2(g)
zinc sulfide + oxygen + Heat → zinc oxide + sulfur dioxide
CuFeS2 + O2 → 2Cu2O + 2Fe2O3
+ SO2
sulfide ore + oxygen → copper oxide + iron oxide + sulfur
dioxide
Calcination:
- Calcination is a process of heating ores or metal carbonate and hydrated compounds in the presence of limited air. This is done to drive off impurities and convert the ore into an oxide.
ZnCO3(s) + heat → ZnO(s) + CO2(g)
Zinc carbonate + heat → zinc oxide + Carbon dioxide
CaCO3(s) + heat → CaO(s) + CO2(g)
Lime stone + heat → Calcium oxide(quick lime)+Carbon
dioxide
Al2O3.2H2O(s) + heat
→ 2Al2O3(s) + 2H2O(l)
2Fe2O3.3H2O(s) + heat
→ 2Fe2O3(s) + 3H2O(l)
Reaction of Metals with Water or steam:
- Metals react with water and form a metal oxide and hydrogen gas. Because Metal oxides are soluble in water, so, dissolve in it to further form metal hydroxide. But all metals do not react with water.
Metal + Water → Metal oxide + Hydrogen
Metal oxide + Water → Metal hydroxide
- Highly reactive metals, such as alkali metals (Group 1, like- sodium, potassium) and alkaline earth metals (Group 2, like- calcium, magnesium) can easily react with water and form metal hydroxides and release hydrogen gas.
- The reaction with water is often vigorous.
- These metals are stored in oil or kerosene to prevent contact with moisture and air.
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
+ Heat
2K(s) + 2H2O(l) → 2KOH(aq) + H2(g)
+ Heat
- Calcium is Moderately Reactive Metals. The reaction of calcium with water is less vigorous. The heat evolved is not sufficient for the hydrogen to catch fire.
Ca(s) + 2H2O(l) (cold)→ Ca(OH)2(aq)
+ H2(g)
In this reaction, Bubbles of hydrogen gas formed stick to the surface of metal so Calcium starts floating.
Mg+2H2O(hot)→Mg(OH)2+H2
- Magnesium reacts with hot water and form magnesium hydroxide and hydrogen. It also starts floating due to the bubbles of hydrogen gas.
- Less reactive metals, like- Aluminum (Al), Zinc (Zn) and Iron (Fe) do not easily react with hot or cold water.
- But they react with steam to form the metal oxide and hydrogen.
2Al(s) + 3H2O(g) (Steam)→ Al2O3(s)
+ 3H2(g)
3Fe(s) + 4H2O(g)(Steam) → Fe3O4(s)
+ 4H2(g)
Zn+H2O(steam)→ZnO+H2
- Lead, copper, silver and gold are metals that do not react with water.
Reaction of Metals with Acids:
Reactions of Metals with Dilute
Hydrochloric Acid(HCl):
- Metals such as magnesium (Mg), aluminum (Al), zinc (Zn) and iron (Fe) react with dilute hydrochloric acid (HCl) to produce metal chloride and hydrogen gas (H2).
- The reactions are as follows:
Magnesium: Mg(s) + 2HCl(aq) → MgCl2(aq)
+ H2(g)
Aluminum: 2Al(s) + 6HCl(aq) → 2AlCl3(aq)
+ 3H2(g)
Zinc: Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
Iron: Fe(s) + 2HCl(aq) → FeCl2(aq) + H2(g)
Reaction with Nitric Acid:
- When metals react with nitric acid (HNO3), with most metals, hydrogen gas is not evolved. This is because nitric acid is a strong oxidizing agent. It oxidizes the hydrogen gas produced to water. Nitric acid itself gets reduced to one of the nitrogen oxides, like- nitrogen dioxide (NO2), nitrogen monoxide (NO), or nitrous oxide (N2O).
- For example, the reaction with magnesium (Mg) can be represented as:
3Mg(s) + 8HNO3(aq) → 3Mg(NO3)2(aq)
+ 4H2O(l) + 2NO(g)
- Only Mg and Mn, react with very dilute nitric acid to liberate hydrogen gas.
Mg(s) + 2HNO3(dilute) → Mg(NO3)2(aq)
+ H2(g)
Mn(s) + 2HNO3(dilute) → Mn(NO3)2(aq)
+ H2(g)
Reactivity of Metals with HCl:
- The reactivity of metals with dilute hydrochloric acid decreases in the order:
- This means that magnesium is the most reactive among the given metals when reacting with HCl.
- Copper does not react with dilute HCl.
Metals reaction with Solutions of other
Metal Salts:
Displacement reaction:
- In a displacement reaction, if a metal A displaces another metal B from its salt solution or molten form, it indicates that metal A is more reactive than metal B.
- The general form of a displacement reaction is:
Metal A + Salt solution of B → Salt solution of A +
Metal B
- This means that metal A effectively replaces metal B in the compound. It’s demonstrating its higher reactivity.
Fe(s) + CuSO4(aq) → FeSO4(aq)
+ Cu(s)
Cu(s) + 2AgNO3(aq) → Cu(NO3)(aq)
+ 2Ag(s)
Mg(s) + CuSO4(aq) → MgSO4(aq)
+ Cu(s)
- These reactions provide clear evidence of how metals compare in terms of reactivity.
The Reactivity Series:
The reactivity or activity series is a
systematic arrangement of commonly used metals in the sequence their decreasing
reactivity or activity.
- Aqua regia is a freshly mix of two strong acids, hydrochloric and nitric acid.
- It's in a 3:1 ratio. It can dissolve gold that neither of these individual acids can do this on their own.
- This highly corrosive and fuming liquid stands out as one of the rare reagents capable of dissolving both gold and platinum.
How Do Metals and Non-Metals React?
- Metals and non-metals have distinct reactions due to their electronic configurations and properties.
- Metals tend to lose electrons from their outermost shell and form positively charged ions (cations) when they react. This happens because metals have few electrons in their outermost shell and they want to achieve a stable, full outer shell.
- Non-metals tend to gain electrons to achieve a full outer shell. They form negatively charged ions (anions) when they react.
|
Types of Element |
Element |
Atomic Number |
Number of electron in shells |
|||
|
K |
L |
M |
N |
|||
|
Noble
gases |
Helium (He) |
2 |
2 |
- |
- |
- |
|
|
Neon
(Ne) |
10 |
2 |
8 |
- |
- |
|
|
Argon (Ar) |
18 |
2 |
8 |
8 |
- |
|
Metals |
Sodium
(Na) |
11 |
2 |
8 |
1 |
- |
|
|
Magnesium (Mg) |
12 |
2 |
8 |
2 |
- |
|
|
Aluminium
(Al) |
13 |
2 |
8 |
3 |
- |
|
|
Potassium (K) |
19 |
2 |
8 |
8 |
1 |
|
|
Calcium
(Ca) |
20 |
2 |
8 |
8 |
2 |
|
Non-metals |
Nitrogen (N) |
7 |
2 |
5 |
- |
- |
|
|
Oxygen
(O) |
8 |
2 |
6 |
- |
- |
|
|
Fluorine (F) |
9 |
2 |
7 |
- |
- |
|
|
Phosphorus
(P) |
15 |
2 |
8 |
5 |
- |
|
|
Sulphur (S) |
16 |
2 |
8 |
6 |
- |
|
|
Chlorine
(Cl) |
17 |
2 |
8 |
7 |
- |
Example-1: Formation of Sodium Chloride -
- Consider the electronic configurations of sodium (Na) and chlorine (Cl). Sodium has one electron in its outermost shell, while chlorine has seven electrons in its outermost shell, requiring one more electron to complete its octet.
- When sodium and chlorine react, sodium loses an electron, forming a sodium cation (Na+), while chlorine gains an electron, forming a chloride anion (Cl—).
- These oppositely charged ions (Na+ and Cl—) are attracted to each other by strong electrostatic forces of attraction, resulting in the formation of the ionic compound sodium chloride (NaCl).
Example-2: Formation of Magnesium Chloride -
- Ionic compounds, like sodium chloride and magnesium chloride, are formed through the transfer of electrons from a metal (cation) to a non-metal (anion).
- In the case of magnesium chloride (MgCl2), magnesium (Mg) forms a cation (Mg2+), while chlorine forms chloride anions (Cl—).
- This process of electron transfer between metals and non-metals leads to the creation of ionic compounds, where strong electrostatic attractions hold the ions together. Ionic compounds play a significant role in chemistry and have various practical applications.
Properties of Ionic compounds:
- Ionic compounds are also referred to as electrovalent compounds, and they consist of aggregates of oppositely charged ions.
- Ionic compounds are electrically neutral compounds composed of positively charged ions (cations) and negatively charged ions (anions).
- Ionic compounds which consist of only two types of elements their nomenclature follow a simple pattern: the name of cation comes first and then write name of the anion.
- These compounds are held together by strong electrostatic attractions between ions of opposite charges.
- Examples : MgCl2, CaO, MgO, and NaCl
- Electrostatic Forces: The primary forces holding ionic compounds
together are electrostatic forces of attraction between positively charged
cations and negatively charged anions. These forces are responsible for the
stability and structure of the compound.
Some general properties of Ionic
compounds are:
- High Melting and Boiling Points:
- Ionic compounds generally have high melting and boiling points. This is because a significant amount of energy is required to overcome the strong electrostatic forces of attraction between oppositely charged ions.
Melting and boiling points of some ionic compounds
Ionic Compound
Melting Point (K)
Boiling Point (K)
In Kelvin (K)
In degrees Celsius (°C)
In Kelvin (K)
In degrees Celsius (°C)
NaCl
1074
801
1686
1413
LiCl
887
614
1600
1327
CaCl2
1045
772
1900
1627
CaO
2850
2577
3120
2847
MgCl2
981
708
1685
1412
- Solubility:
- Most of the ionic compounds are soluble in water because water molecules break them apart due to polar nature of water. They are insoluble in solvents like kerosene, Petrol etc. However, not all ionic compounds are equally soluble. For example:
- NaCl forms a three-dimensional crystal structure composed of Na+ (sodium) and Cl− (chloride) ions.
- Within this crystal, the ions are held together by strong electrostatic forces of attraction due to their opposite charges.
- When a crystal of NaCl is placed in water, the interaction begins between the ions and water molecules.
- The partially positively charged ends of water molecules are attracted to the Cl− ions, while the negatively charged ends are attracted to the Na+ ions.
- This ion-dipole interaction between ions and water molecules plays a crucial role in breaking the strong electrostatic forces within the crystal lattice.
- As a result of this interaction, the crystal gradually dissolves or becomes soluble in the water.
- This process illustrates how the interaction between ions and water molecules enables the dissolution or solubility of ionic compounds like NaCl in a liquid medium.
- Conductivity:
- In the solid state, ionic compounds do not conduct electricity because the ions are held in a fixed lattice and cannot move. In the molten state (when heated to high temperatures) or when dissolved in water, they conduct electricity because the ions are free to move and carry electric charge.
- For example: Such ionic compounds like NaCl do not conduct electricity in solid state, but when dissolved in water or in a molten state, they will conduct electricity.
- Physical nature:
- Ionic compounds have a regular, repeating crystal lattice structure in the solid state. The arrangement of cations and anions is highly ordered and three-dimensional.
- Ionic compounds are brittle due to the arrangement of ions in a crystal lattice structure.
- Many ionic compounds in solid state are transparent or translucent, allowing light to pass through.
- Ionic compounds are non-volatile at room temperature, meaning they do not easily evaporate into a gas.
- Some ionic compounds are hygroscopic nature. Its mean they easily absorb water vapor from the environment and may become hydrated when exposed to humid conditions.
Occurrence of Metals:
Sources of Metals
- The Earth's crust is the primary source of metals. This solid outer layer of the Earth contains a wide variety of metallic elements and compounds.
- Seawater is another source of metals, as it contains soluble salts like sodium chloride and magnesium chloride.
- The occurrence of metals can be broadly categorized into three main types:
- Native metals: Native metals are metals that exist in nature in
their pure, uncombined form like Au, Pt.
- Minerals: Elements or compounds that occur naturally in the
Earth's crust are known as minerals. These minerals can be composed of various
elements or compounds.
- Ores: In some geological locations, minerals contain a
high concentration of a particular metal which makes it feasible to extract and
economically viable. These special minerals with a high metal content are
referred to as ores. Ores are important because they provide a
concentrated and accessible source of metals for industrial use.
Extraction of Metals:
Metals are extracted from their ores
using various techniques. The choice of technique depends on the reactivity of
the metal and the nature of its ore. Following are the extraction process and
categorization of metals based on reactivity:
- Reactivity Series:
- Metals vary in their reactivity and it is depends on, how they are found in nature.
- The least reactive metals are found in their pure (free) state in nature. Examples- gold, silver, platinum and copper.
- Some moderately reactive metals are found as compounds in ores, mainly as oxides, sulphides or carbonates. Examples- zinc, iron and lead.
- The highly reactive metals, such as potassium (K), sodium (Na), calcium (Ca), magnesium (Mg) and aluminum (Al) are never found in nature as free elements.
K
Na
Ca
Mg
Al
Electrolysis
(High Reactivity)
never found as free elements
Zn
Fe
Pb
Cu
Reduction using carbon (Medium reactivity) found as oxides, sulphides or carbonates.
Ag
Au
Found in Native state or pure state
(low reactivity) found in native state
- Categorization by Reactivity:
- Metals are categorized into three groups based on reactivity:(i) low reactivity (ii) medium reactivity (iii) high reactivity.
- Low-reactivity metals include those found in their pure state in nature.
- Medium-reactivity metals are mainly found in the form of oxides, sulphides or carbonates.
- High-reactivity metals are never found in nature as free elements.
Extraction Steps:
The extraction of pure metals from ores
involves multiple steps. These steps are summarized as follows:
- Crushing and grinding of the ore to reduce its size.
- Concentration of the ore to remove impurities.
- Reduction of the ore to obtain the metal using suitable reducing agents.
- Purification of the obtained metal through refining processes.
Different techniques are applied for
extracting metals based on their reactivity and the characteristics of their
ores.
Enrichment of Ores:
- Ores extracted from the Earth often contain significant impurities, such as soil and sand, collectively referred to as gangue.
- To obtain pure metals from these ores, the impurities must be separated or removed before the extraction process.
- Enrichment of ores is a preliminary step in the extraction of metals to ensures that the final product is as pure as possible.
- The techniques applied for this purpose will vary depending on the nature of the ore and the specific impurities present.
Extracting Metals Low in the Activity
Series:
- Metals lows in reactivity series are very un-reactive.
- Their oxides can be reduced to obtain the pure metal by heating alone.
- The reduction of metal oxides to obtain pure metals is a important step in the extraction process.
Example (i)- Mercury (Hg):
- Cinnabar (HgS) is an ore of mercury.
- When heated in the presence of air, it undergoes a two-step process.
(i) It converts into mercuric oxide (HgO):
2HgS(s) + 3O2(g) + Heat → 2HgO(s) + 2SO2(g)
(ii) Mercuric oxide (HgO) is further heated to yield mercury (Hg) and oxygen gas (O2):
2HgO(s) + Heat → 2Hg(l) + O2(g)
Example(ii) - Copper (Cu):
- Copper sulfide (Cu2S) is an ore of copper.
- When heated in the presence of air, it undergoes a two-step process.
(i) It converts copper sulfide (Cu2S) to copper oxide (Cu2O):
2Cu2S + 3O2(g) + Heat → 2Cu2O(s)
+ 2SO2(g)
(ii) Copper oxide (Cu2O) is further heated to yield copper (Cu) and SO2 gas:
2Cu2O + Cu2S + Heat → 6Cu(s)
+ SO2(g)
Example (iii) – Lead (Pb):
- Galena (PbS) is an ore of lead.
- When heated in the presence of air, it undergoes a two-step process.
(i) It converts PbS to PbO:
2PbS + 3O2(g) + Heat → 2PbO(s)
+ 2SO2(g)
(ii) PbO is further heated to yield Lead (Pb) and SO2 gas:
PbS(s) + 2PbO(s) → 2Pb(s) + SO2(g)
Extraction of Metals in the Middle of
the Activity Series:
- Metals in the middle of the reactivity series, such as iron, zinc, lead, and copper, are moderately reactive. They are commonly found in nature as sulphides or carbonates.
- It is easier to extract these metals from their oxides compared to their sulphides and carbonates. Therefore, the sulphide and carbonate ores must first be converted into metal oxides. This process called Oxidation.
- This conversion is achieved by heating the sulphide ores strongly in the presence of excess air, a process known as "Roasting".
- Carbonate ores are converted into oxides by strong heating in limited air, which is called "calcination".
- Example: The chemical reactions during the roasting and calcination of zinc ores can be represented as follows:
- Roasting:
2ZnS(s) + 3O2(g) + Heat → 2ZnO(s) + 2SO2(g)
- Calcination:
ZnCO3(s) + Heat → ZnO(s) + CO2(g)
- After converting the ores to metal oxides, the next step is to reduce these oxides to obtain the corresponding metals in presence of reducing agent like carbon. This is called smelting.
- Example: When zinc oxide is heated with carbon, it is reduced to produce metallic zinc:
ZnO(s) + C(s) → Zn(s) + CO(g)
Displacement reactions:
- By using carbon (coke) as reducing agent, displacement reactions can use for this purpose
- Highly reactive metals like sodium, calcium, and aluminum can serve as powerful reducing agents because they can displace metals of lower reactivity from their compounds.
Example:
- When manganese dioxide (MnO2) is heated with aluminum powder (Al), a displacement reaction occurs, leading to the production of molten manganese and aluminum oxide:
3MnO2(s) + 4Al(s) → 3Mn(l) + 2Al2O3(s)
+ Heat
Oxidation and Reduction in Displacement Reactions:
- In displacement reactions, some substances get oxidized and others get reduced and it is highly exothermic, releasing a significant amount of heat that the metals are produced in the molten state. This is known as Thermit reactions.
- Example: The reaction of iron(III) oxide (Fe2O3) with aluminum, the following reaction takes place:
Fe2O3(s) + 2Al(s) → 2Fe(l) +
Al2O3(s) + Heat
- Above reaction is used in application such as joining railway tracks or repairing cracked machine parts.
Extraction of Metals towards the Top of
the Activity Series:
- Metals high up in the reactivity series are very reactive. For example: sodium, magnesium, calcium and aluminum.
- These highly reactive metals cannot be obtained from their compounds by heating with carbon.
- Carbon is ineffective in reducing the oxides of these metals because they have a higher affinity for oxygen than carbon.
- To extract these metals, electrolytic reduction is employed.
Electrolytic Reduction:
- Sodium, magnesium, and calcium are obtained by the process of electrolysis of their molten chlorides.
- In this process:
- The metals are deposited at the cathode, which is the negatively charged electrode.
- Gasses like oxygen Chlorine, is liberated at the anode, which is the positively charged electrode.
- The soluble impurities go into the solution and insoluble impurities settle down at the bottom of the anode and are known as anode mud.
Example 1: Electrolytic refining of sodium:
At the cathode (reduction) : Na⁺ + e⁻ → Na(s) (sodium
metal deposited)
At the anode (oxidation) : 2Cl⁻ → Cl₂ + 2e⁻ (chlorine
gas liberated)
Example 2: Electrolytic refining of Aluminium:
At the cathode (reduction) : 2Al3+ + 6e⁻
→ 2Al(s) (Aluminium metal deposited)
At the anode (oxidation) : 6O2⁻ → 3O2
+ 12e⁻ (Oxygen gas liberated)
Corrosion and Prevention:
- Silver articles become black when exposed to air due to a reaction with sulphur, forming a coating of silver sulphide.
2Ag(s) + H2S (from air) → Ag2S
(black) + H2(g)
- Copper reacts with moist carbon dioxide in the air, leading to a loss of its shine and the formation of green basic copper carbonate.
Cu(s) + H2O (moisture) + CO2 (from
air) → CuCO3.Cu(OH)2 (green)
- Iron exposed to moist air over time acquires a brown, flaky substance called rust.
4Fe(s) + 3O2 (from air) + H2O
(moisture) → 2Fe2O3.H2O (rust)
Prevention of Corrosion:
- Methods to prevent iron rusting include painting, oiling, greasing, chrome plating and making alloys.
- Galvanization: In this process a thin layer of molten zinc apply on iron articles. Zinc provides protection from corrosion.
- Electroplating: In this process, a thin coating of one metal with another by using electric current. Example: Nickel painting, silver painting.
- Sacrificial protection: When Magnesium coating on articles of iron or steel, it act as cathode undergoes a reaction (sacrifice) instead of iron and protects the articles because Magnesium is more reactive than iron.
Alloying and Its Benefits:
- Alloying improves the properties of metals.
- Pure iron is soft and easily stretches when hot. Mixing with a small amount of carbon (about 0.05%) makes it hard and strong.
- Iron mixed with nickel and chromium results in stainless steel, which is both hard and corrosion-resistant.
- Alloys can be created by mixing metals or metals with non-metals, offering a way to alter the properties of the base metal.
- Alloys are homogeneous mixtures prepared by melting the primary metal and dissolving other elements in specific proportions.
- Amalgam: If mercury is one of the metals, the alloy is
known as an amalgam.
- Alloys have lower electrical conductivity and melting points compared to pure metals.
- Examples of alloys: brass (copper and zinc), bronze (copper and tin), and solder (lead and tin).
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