Group Properties and Elements

Ans. A catalytic converter is a device used in the exhaust of an automobile which converts more harmful gases produced in the engine to such gases which do not pollute the atmosphere. Platinum, palladium and rhodium are the catalysts used in catalytic converters. e.g: They convert CO to CO₂ and oxides of nitrogen NOₓ to N₂ before it enters in air.

Ans. The trend in chemical reactivity for alkali metals and halogens can be explained by looking at their atomic size and how it affects their ability to lose or gain electrons.

Alkali Metals: All the elements present in group 1 have ns¹ configuration in their outermost shells. This single electron can be removed easily which makes these metals very reactive except the first element hydrogen which is a gas and a non-metal. When we move from top to bottom in this group, the atomic size increases and nuclear charge decreases. Owing to this, it becomes easier for elements to lose electron down the group which increases reactivities of the lower members of the group.

Reaction with water: Lithium reacts with water steadily giving hydrogen and lithium hydroxide. Sodium reacts vigorously while potassium reacts violently with water giving their respective water soluble hydroxides.

Li_(s) + H₂O_(ℓ) —Heat→ LiOH_(aq) + H₂_(g)
Na_(s) + H₂O_(ℓ) —Heat→ NaOH_(aq) + H₂_(g)
K_(s) + H₂O_(ℓ) —Heat→ KOH_(aq) + H₂_(g)

Reaction with halogens: Reaction of these metals with chlorine becomes more vigorous as we go down the group.

2Li_(s) + Cl₂_(g) —Heat→ 2LiCl_(s)
2Na_(s) + Cl₂_(g) —Heat, Vigorous reaction→ 2NaCl_(s)
2K_(s) + Cl₂_(g) —Heat, Violent reaction→ 2KCl_(s)

Halogens: Halogens react with alkali and alkaline earth metals to give salts. These elements are thus named was halogens which means salt-forming elements. Unlike metals, the reactivity of halogens decreases from top to bottom in the group. This is due to the fact that atomic size increases down the group and tendency to accept electron from other atoms decreases making them less reactive. e.g. Reactivity order of halogens F >Cl> Br >I.

Ans. Metals are generally tough and strong due to their unique atomic structure and bonding characteristics, while non-metals do not exhibit these properties for several reasons.

Metallic Bonding: Metals have metallic bonds, which involve a sea of electrons that are free to move throughout the metal lattice. This allows metals to absorb and distribute energy effectively, making them tough and able to withstand significant stress without breaking. The strong attraction between the positively charged metal ions and the electrons contributes to their strength.

Crystal Structure: Metals typically have a crystalline structure which means their atoms are arranged in a regular, repeating pattern. This arrangement allows for the layers of atoms to slide over each other without breaking the metallic bond contributing to their ductility and toughness.

In contrast, non-metals have different bonding characteristics:

Covalent Bonding: Non-metals usually form covalent bonds which are generally weaker than metallic bonds. This results in non-metals being more brittle and less able to withstand stress.

Molecular Structure: Many non-metals exist as discrete molecules (like gases or simple covalent compounds), which do not have the same structural integrity as metals. This lack of a strong, organized structure contributes to their lower toughness and strength.

Ans. Alkali metals and halogens are indeed very reactive elements, but they have opposite roles in chemical reactions due to their positions in the periodic table and their electronic configurations.

Alkali Metals:
Alkali metals are located in Group 1 of the periodic table. They have one electron in their outermost shell which they readily lose to achieve a stable electronic configuration.

This tendency to lose one electron makes alkali metals highly reactive, especially with non-metals. When they react, they form positive ions (cations) with a charge of +1.

For example, when sodium (Na) reacts with chlorine (Cl), sodium loses one electron to form Na⁺ ions.

Halogens:
Halogens are located in Group 17 of the periodic table. They have seven electrons in their outermost shell and need one more electron to achieve a full outer shell which gives them stability.

This makes halogens highly reactive as well but they tend to gain an electron during reactions. When they react, they form negative ions (anions) with a charge of -1.

For example, chlorine (Cl) gain one electron from sodium to form Cl⁻ ions.

Ans. Bond length between hydrogen and halogen increases down the group because as the halogen atom gets bigger the bonding pairs of electron get further away from the halogen nucleus. The bond between hydrogen and halogen therefore gets weaker. The weaker the bond, the less heat energy it will need to break it. Hence the thermal stability of hydrogen halides decreases down the group.

Bond Strength:
The H-Cl bond is stronger than the H-Br bond. This is because atomic size of chlorine is smaller than bromine. A stronger bond is less likely to break under thermal conditions.

The bond dissociation energy for H-Cl is about 431 kJ/mol, while for H-Br, it is around 366 kJ/mol. The lower bond strength in HBr makes it easier to break apart when heated.

Molecular Size and Stability:
Bromine is larger than chlorine, which leads to a longer and weaker bond in HBr. The increased distance between the hydrogen and bromine atoms results in less effective orbital overlap and a more unstable bond.

Thermal Decomposition:
When subjected to heat, HBr can decompose into hydrogen (H₂) and bromine (Br₂) more readily than HCl to decompose into hydrogen (H₂) and chlorine (Cl₂). This decomposition reaction is more favorable for HBr due to its weaker bond.

2 HBr_(g) ⟶ H₂_(g) + Br₂_(g)
2 HCl_(g) ⟶ H₂_(g) + Cl₂_(g)

Ans.

Ans. Platinum was originally used as a catalyst in the contact process for the manufacture of sulphuric acid. Because it is expensive catalyst, however, rendered inactive due to the presence of arsenic as impurity in sulphur dioxide. That is why vanadium pentoxide (V₂O₅) is now preferred as a catalyst.

2SO₂_(g) + O₂_(g) —V₂O₅→ 2SO₃_(g)
SO₃_(g) + H₂O_(ℓ) ⟶ H₂SO₄_(aq)

Ans. Sodium plays a crucial role as a heat transfer agent in atomic nuclear power plants, particularly in sodium-cooled fast reactors. The primary property of sodium that is utilized in this role is its excellent thermal conductivity and its ability to remain liquid at high temperature.

High Thermal Conductivity:
Sodium has a high thermal conductivity, which means it can efficiently transfer heat from the reactor core to the heat exchanger. This property is essential in maintaining the temperature within the reactor and ensuring that the heat generated from nuclear fission is effectively removed.

Liquid State at High Temperatures:
Sodium remains in a liquid state at elevated temperatures (its melting point is around 98 °C). This allows it to circulate freely within the reactor system, absorbing heat from the reactor core and transferring it to the secondary cooling system.

Low Neutron Absorption:
Sodium has a low neutron absorption cross-section, which means it does not interfere with the nuclear fission process. This property is vital in fast reactors where maintaining the neutron economy is crucial for sustaining the fission reaction.

Chemical Stability:
Sodium is chemically stable and can be used in a closed-loop system without significant corrosion or degradation which is important for the durability and safety of the reactor system.

Ans. Lithium behaves differently from the other alkali metals due to its unique properties and smaller size.

Size and Charge Density:
Lithium is the smallest alkali metal, which gives it a higher charge density as compared to the larger alkali metals like sodium, potassium, and others. This high charge density leads to stronger interactions with surrounding ions or molecules, resulting in different chemical behavior.

Ionization Energy:
Lithium has high ionization energy than the other alkali metals. This means it requires more energy to remove the outermost electron from lithium compared to the other alkali metals. As a result, lithium is less reactive than other alkali metals.

Covalent Bonding:
Lithium tends to form covalent bonds more readily than the other alkali metals, which typically form ionic bonds. This is due to the small size and high charge density of lithium which allows it to share electrons more effectively.

Solubility in Organic Solvents:
Unlike other alkali metals, lithium can dissolve in some organic solvents. This property makes lithium useful in organic synthesis and various chemical reactions.

Lithium Compounds:
The compounds formed by lithium, such as lithium carbonate or lithium hydroxide have different properties compared to similar compounds of other alkali metals. For example, lithium carbonate is less soluble in water than sodium carbonate.

Ans. Aluminium is commonly used in the manufacture of cooking utensils for several reasons, while magnesium is not typically chosen for this purpose. Such as

Corrosion Resistance:
Aluminium has a natural oxide layer that protects it from corrosion. This makes it ideal for cooking utensils that are often exposed to moisture and acidic foods. Magnesium, on the other hand, is more chance to corrode which can lead to deterioration over time.

Weight and Strength:
Aluminium is lightweight yet strong, making it easy to handle while still being durable enough for cooking applications. Magnesium is even lighter than aluminium, but it is also more brittle and can break or deform under stress, making it less suitable for cookware.

Thermal Conductivity:
Aluminium has excellent thermal conductivity, allowing for even heat distribution when cooking. This property is essential for effective cooking. Although magnesium also has good thermal conductivity, its other drawbacks make it less desirable for cooking utensils.

Cost and Availability:
Aluminium is relatively inexpensive and widely available, making it a practical choice for mass production of cooking utensils. Magnesium can be more costly and less readily available for such applications.

Reactivity:
Magnesium is more reactive than aluminium, especially at higher temperatures. This reactivity can lead to undesirable chemical reactions with certain foods or cooking processes, making it less suitable for cookware.

Ans. Introduction of Alkali Metals
Elements present in a group of the periodic table show similar chemical properties owing to the presence of same number of electrons in their outermost shells. However, a small variation in the chemical properties of elements is expected because the atomic size increases down the group.

Definition:
All the elements present in group 1 have ns¹ configuration in their outer shells. They are also called alkali metals.

Reactivity:
This single electron can be removed easily which makes these metals very reactive except the first element hydrogen which is a gas and a non-metal. When we move from top to bottom in this group the atomic size increases. Owing to this, it becomes easier for elements to lose electron down the group which is reflected in the increased reactivties of the lower members of the group.

Reaction with water
Lithium reacts with water steadily giving hydrogen and lithium hydroxide. Sodium reacts vigorously while potassium reacts violently with water giving their respective water soluble hydroxides.

Li_(s) + H₂O_(ℓ) ⟶ LiOH_(aq) + H₂_(g)
Na_(s) + H₂O_(ℓ) ⟶ NaOH_(aq) + H₂_(g)
K_(s) + H₂O_(ℓ) ⟶ KOH_(aq) + H₂_(g)

Reaction with halogens:
Reaction of these metals with chlorine becomes more vigorous as we go down the group.

2Li_(s) + Cl₂_(g) —Heat→ 2LiCl_(s)
2Na_(s) + Cl₂_(g) —Heat, Vigorous reaction→ 2NaCl_(s)
2Ka_(s) + Cl₂_(g) —Heat, Violent reaction→ 2KCl_(s)

Ans.

• Halogens generally do not act as reducing agents in their elements form (e.g F₂,Cl₂, Br₂,I₂) because they tend to gain electrons than lose them.
• Instead, Halogens are typically strong oxidizing agents However their halide Ions (e. g F⁻, Cl⁻, Br⁻,I⁻) can act as reducing agents under specific conditions.
• The reducing strength of halides increase in the order F⁻< Cl⁻ < Br⁻<I⁻
• Halogens are reducing agents & their reducing power decreases down the group.

F₂ > Cl₂ > Br₂ > I₂

This fact gives a unique property to halogens when a halogen having more oxidizing reducing power displace an ion of another halogen from its compound.

Cl₂_(g) + 2NaBr_(aq) ⟶ 2NaCl_(aq) + Br₂_(g)
Cl₂_(g) + 2NaI_(aq) ⟶ 2NaCl_(aq) + I₂_(v)
Br₂_(aq) + 2NaI_(aq) ⟶ 2NaBr_(aq) + I₂_(v)

Ans: Halogens react with hydrogen to give hydrogen halides, Hydrogen halides behave as strong acids in water

H₂ + X₂ ⟶ 2HX

All halides exist in gaseous state at ordinary temperature except hydrogen fluoride which is a liquid. Bond length between hydrogen and halogen increases down the group because as the halogen atom gets bigger the bonding pairs of electron get further away from the halogen nucleus. The bond between hydrogen and halogen therefore gets weaker. The weaker the bond, the less heat energy it will need to break it. Hence the thermal stability of hydrogen halides decreases down the group.

HF > HCl> HBr > HI

Ans. Elements present in group 18 of the modern periodic table are called Noble elements

Properties of Noble gases:

i. All Noble elements are monoatomic gases having very low boiling points Helium (He), Neon (Ne), Krypton (Kr), Xenon (Xe) and Radon (Rn).
ii. All these gases have eight electrons (s²p⁶) in their outermost shells except He which has s² electronic configuration.
iii. Since their outer shells are complete. They show very little chemical reactivity.

Ans. Definition of metals:

Metals are defined as the elements which can generally form cations easily. They also tend to form metallic bond.

Physical properties of metals:

i. Metals can be hammered into thin sheets. This property is called malleability.
ii. Metals can also be drawn into wires and this property is named as ductility
iii. Metallic bond in metals allows metals to be the best conductor of heat and electricity.
iv. Metals are lustrous which means that they have a shiny appearance.
v. Due to high tensile strength metals can hold heavy weights.
vi. When metals are hit by an object, they make a ringing sound.
vii. Metals cannot be cut easily because they are hard substances
viii. Due to the presence of strong metallic bond metals generally have high melting and boiling points.
ix. Their densities are also very high.
x. Alkali metals being soft metals are treated as exceptions

Example: Metals include copper, silver, iron, lead aluminum, gold, platinum, zinc etc.

Ans. Physical properties of non-metals are given below:

i. Non-metals show a greater variety of colours and physical states compared to metals.
ii. Non-metals cannot be beaten into thin sheets because being brittle they break into pieces when hammered. Sulphur and phosphorous exist in powdered forms and cannot be made into sheets: Non-metals cannot be melted and drawn into wires.
iii. Non-metals do not have free electrons due to which the bonds between their atoms are weak and they break down when stretched.
iv. As there are no free electrons so non-metals cannot conduct heat and electricity. Graphite is the only exception. It conducts electricity because of its special crystalline arrangements.
v. Non-metals cannot be polished because they either exist in powder or gaseous form. Most of the powders are dull in texture.
vi. Due to non-ductile and non-malleable properties, non-metals are not strong at all. Their bonds being weak break easily.
vii. All non-metals have low melting and boiling points. The melting point of sulphur is 115°C. Graphite and diamond have high melting points and these are exceptions.
viii. Non-metals have low densities as compared to metals. This means that in non-metals atoms are not strongly bound with each other.

Examples: Non-metals are oxygen, nitrogen, chlorine, sulphur, carbon, bromine, etc.