Day: September 15, 2024

Periodic Classification of Elements

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Periodic classification of elements and periodicity (Solved exercise PTB)

Q4:
What are the improvements made in Mendeleev’s periodic table?

  • Mendeleev left gaps for undiscovered elements and predicted their properties accurately.
  • He arranged elements based on increasing atomic mass, but later corrections were made based on atomic number.
  • Grouping of elements with similar chemical properties was better explained by the periodic table developed by Moseley using atomic number.

Q5:
How does the classification of elements in different blocks help in understanding their chemistry?

  • Elements are classified into s-block, p-block, d-block, and f-block based on their valence electron configurations, which helps in understanding their chemical behavior.
  • s-block elements are typically metals, p-block elements include non-metals, metalloids, and some metals, d-block elements are transition metals, and f-block contains lanthanides and actinides with complex chemistry.

Q6:
How do you justify the position of hydrogen at the top of various groups?

  • Hydrogen is placed at the top of Group 1 due to its ability to form a +1 oxidation state like alkali metals, but it can also form a -1 oxidation state (like halogens in Group 17).
  • Its properties resemble both alkali metals (in forming H⁺ ions) and halogens (in forming H⁻ ions), which justifies its unique position.

Q7:
Why is the ionic radii of negative ions larger than the size of their parent atoms?

  • When an atom gains an electron to form a negative ion (anion), the repulsion between the electrons increases, causing the electron cloud to expand. This results in the anion having a larger radius than the parent atom.

Q8:
Why does ionization energy decrease down the group and increase along a period?

  • Down a group: Ionization energy decreases because the atomic size increases, and the outermost electrons are farther from the nucleus, reducing the attraction.
  • Across a period: Ionization energy increases because atomic size decreases and nuclear charge increases, making it harder to remove an electron.

Q9:
Why is the second value of electron affinity of an element usually shown with a positive sign?

  • The second electron affinity is positive because after the addition of one electron, the atom becomes negatively charged. Adding another electron to a negative ion requires energy to overcome the repulsion between the added electron and the negative ion.

Q10:
Why does metallic character increase from top to bottom in a group of metals?

  • As you move down a group, the atomic size increases, and the outermost electrons are farther from the nucleus, making them easier to lose. This enhances the metallic character, which is associated with the ease of losing electrons.

Q11:
Explain the variation in melting points along the short periods.

  • Melting points tend to increase across a period until reaching Group 14, where they reach a maximum and then decrease. This is due to the increasing strength of metallic bonding or covalent bonding in elements like silicon and then weaker van der Waals forces in non-metals.

Q12:
Why is the oxidation state of noble gases usually zero?

  • Noble gases have a completely filled valence shell, making them highly stable and inert. They do not easily lose or gain electrons, which is why their oxidation state is generally zero.

Q13:
Why is diamond a non-conductor and graphite a good conductor?

  • Diamond: In diamond, each carbon atom is tetrahedrally bonded to four other carbon atoms, and all valence electrons are used in bonding. There are no free electrons to conduct electricity.
  • Graphite: In graphite, each carbon atom is bonded to three other carbon atoms, with one delocalized electron per atom. These free electrons can move through the layers, allowing graphite to conduct electricity.

Q14:
Give a brief reason for the following:

(a) d- and f-block elements are called transition elements.

  • They are called transition elements because they exhibit properties that are transitional between s- and p-block elements. They have partially filled d- or f-orbitals, which gives them unique chemical and physical properties.

(b) Lanthanide contraction controls the atomic sizes of elements of the 6th and 7th periods.

  • Lanthanide contraction refers to the steady decrease in atomic and ionic sizes of the lanthanides as atomic number increases. This contraction influences the sizes of elements in the 6th and 7th periods, particularly in d- and f-block elements.

(c) The melting and boiling points of elements increase from left to the middle of the s- and p-block elements and decrease onward.

  • From left to the middle of the s- and p-block, elements exhibit stronger metallic or covalent bonding, resulting in higher melting and boiling points. After the middle, the elements become non-metals with weaker van der Waals forces, leading to lower melting and boiling points.

(d) The oxidation states vary in a period but remain almost constant in a group.

  • Across a period, elements can exhibit multiple oxidation states due to the availability of different numbers of valence electrons for bonding. Within a group, the number of valence electrons remains the same, so the oxidation states are more consistent.

(e) The hydration energies of the ions are in the following order: Al³⁺ > Mg²⁺ > Na⁺.

  • Hydration energy is higher for ions with a greater charge and smaller size. Al³⁺ has the smallest size and highest charge, followed by Mg²⁺, and then Na⁺, leading to the given order.

(f) Ionic character of halides decreases from left to right in a period.

  • As you move across a period, the electronegativity of the elements increases. The greater the electronegativity difference between the metal and the halide, the higher the ionic character. This difference decreases across a period, reducing the ionic character.

(g) Alkali metals give ionic hydrides.

  • Alkali metals react with hydrogen to form ionic hydrides (e.g., NaH) where the hydrogen is present as a hydride ion (H⁻) due to the large difference in electronegativity between alkali metals and hydrogen.

(h) Although sodium and phosphorus are present in the same period of the periodic table, their oxides are different in nature: Na₂O is basic while P₄O₁₀ is acidic.

  • Sodium, being a metal, forms a basic oxide (Na₂O), which reacts with water to produce a strong base (NaOH). Phosphorus, being a non-metal, forms an acidic oxide (P₄O₁₀), which reacts with water to form phosphoric acid (H₃PO₄).

s block elements- Solve Exercise (Punjab Boards)

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Q4 (a):
Names, electronic configurations, and occurrence of s-block elements:

  • Elements: The s-block elements include Group 1 (alkali metals: Li, Na, K, Rb, Cs, Fr) and Group 2 (alkaline earth metals: Be, Mg, Ca, Sr, Ba, Ra).
  • Electronic configurations:
  • Alkali metals: General configuration is (ns^1).
    • Li: (1s2 2s1), Na: (1s^2 2s^2 2p^6 3s^1), etc.
  • Alkaline earth metals: General configuration is (ns^2).
    • Be: (1s^2 2s^2), Mg: (1s^2 2s^2 2p^6 3s^2), etc.
  • Occurrence: These elements are found in the Earth’s crust in minerals, salts, and ores (e.g., NaCl, CaCO₃, MgCO₃).

(b):
Peculiar behavior of lithium with respect to other alkali metals:

  • Lithium has smaller atomic and ionic size compared to other alkali metals.
  • It forms a stable oxide (Li₂O) and reacts with nitrogen to form lithium nitride (Li₃N).
  • Lithium salts have more covalent character due to its higher polarization ability.
  • It is less reactive with water than other alkali metals.
  • Shows diagonal relationship with magnesium, resembling its properties more.

Q5:
Trends in the chemical properties of oxides, hydroxides, carbonates, nitrates, and sulfates of Group IA (alkali metals) and Group IIA (alkaline earth metals):

  • Oxides: Alkali metal oxides (e.g., Li₂O) are basic, while alkaline earth metal oxides (e.g., MgO, CaO) are also basic but less soluble.
  • Hydroxides: Alkali metal hydroxides (e.g., NaOH) are strong bases, while alkaline earth metal hydroxides (e.g., Ca(OH)₂) are less soluble and weaker bases.
  • Carbonates: Alkali metal carbonates (e.g., Na₂CO₃) are soluble in water, while alkaline earth metal carbonates (e.g., CaCO₃) are insoluble.
  • Nitrates: Both alkali and alkaline earth nitrates (e.g., NaNO₃, Ca(NO₃)₂) decompose on heating to give oxides, but alkaline earth nitrates decompose more readily.
  • Sulfates: Alkali metal sulfates (e.g., Na₂SO₄) are soluble in water, while alkaline earth metal sulfates (e.g., BaSO₄) are less soluble.

Q6:
Comparison of chemical behavior of lithium with magnesium:

  • Both lithium and magnesium form oxides (Li₂O and MgO) and hydroxides (LiOH and Mg(OH)₂).
  • They form nitrides when reacted with nitrogen (Li₃N, Mg₃N₂).
  • Lithium and magnesium carbonates decompose on heating to form oxides and CO₂.
  • Both do not form peroxides or superoxides, unlike other members of their groups.
  • Both show covalent bonding due to their smaller size and higher ionization energies.

Q7 (a):
Properties of beryllium that make it different from its family members:

  • Beryllium does not react with water, unlike other alkaline earth metals.
  • Its salts are generally covalent rather than ionic.
  • Beryllium forms complexes, e.g., [BeF₄]²⁻.
  • It has a high melting point and does not show characteristic flame coloration.
  • Beryllium oxide (BeO) is amphoteric, unlike other Group II metal oxides, which are basic.

Q8 (b):
Why is an aqueous solution of Na₂CO₃ alkaline?

  • Na₂CO₃ is a salt of a strong base (NaOH) and a weak acid (H₂CO₃). In water, it hydrolyzes to produce OH⁻ ions, making the solution alkaline:
    [ Na₂CO₃ + H₂O → 2Na⁺ + CO₃^{2-} ]
    [ CO₃^{2-} + H₂O → HCO₃^- + OH⁻ ]

Q9 (a):
Describe with a diagram the manufacture of sodium by Down’s cell:

  • Down’s cell is an electrolytic process used to manufacture sodium by the electrolysis of molten NaCl.
  • The cathode is made of iron and the anode is made of graphite.
  • Na⁺ ions are reduced at the cathode to form sodium metal, and Cl⁻ ions are oxidized at the anode to form chlorine gas:
    [ 2NaCl → 2Na + Cl₂ ]

(b):
Three advantages of the Down’s process:

  1. High-purity sodium is obtained.
  2. It is an economical and continuous process.
  3. Chlorine gas, a useful byproduct, is obtained.

Q10:
Compare the physical and chemical properties of alkali metals with those of alkaline earth metals:

  • Physical properties:
  • Alkali metals are softer and have lower melting and boiling points than alkaline earth metals.
  • Alkali metals are more reactive than alkaline earth metals.
  • Alkali metals have lower densities compared to alkaline earth metals.
  • Chemical properties:
  • Alkali metals form monovalent cations (M⁺), whereas alkaline earth metals form divalent cations (M²⁺).
  • Alkali metals react more vigorously with water than alkaline earth metals.
  • Alkali metal oxides are more basic than alkaline earth metal oxides.

Here are the answers to the given questions based on the image:

Q9 (b):
What happens when:

  1. Lithium carbonate is heated?
  • Lithium carbonate (Li₂CO₃) decomposes upon heating to form lithium oxide (Li₂O) and carbon dioxide (CO₂):
    [ Li₂CO₃ → heat Li₂O + CO₂ ]
  1. Lithium hydroxide is heated to red heat?
  • Lithium hydroxide (LiOH) decomposes to form lithium oxide (Li₂O) and water (H₂O):
    [ 2LiOH → Li₂O + H₂O ]
  1. Beryllium is treated with sodium hydroxide?
  • Beryllium reacts with sodium hydroxide (NaOH) to form sodium beryllate (Na₂BeO₂) and hydrogen gas (H₂):
    [ Be + 2NaOH + 2H₂O → Na₂BeO₂ + H₂ ]
  1. Lithium hydride is treated with water?
  • Lithium hydride (LiH) reacts with water to form lithium hydroxide (LiOH) and hydrogen gas (H₂):
    [ LiH + H₂O→ LiOH + H₂ ]

Q10:
Give the formulas of the following minerals:

  1. Dolomite:
    [ \text{CaMg(CO₃)₂} ]
  2. Asbestos:
    [ \text{Mg₃Si₂O₅(OH)₄} ] (Chrysotile, the most common form)
  3. Halite:
    [ \text{NaCl} ] (Sodium chloride)
  4. Natron:
    [ \text{Na₂CO₃·10H₂O} ] (Sodium carbonate decahydrate)
  5. Beryl:
    [ \text{Be₃Al₂Si₆O₁₈} ] (Beryllium aluminum cyclosilicate)
  6. Sylvite:
    [ \text{KCl} ] (Potassium chloride)
  7. Phosphorite:
    [ \text{Ca₅(PO₄)₃F} ] or [ \text{Ca₃(PO₄)₂} ] (Fluorapatite or phosphate rock)
  8. Chile saltpeter:
    [ \text{NaNO₃} ] (Sodium nitrate)

Q11:
Answer the following questions briefly:

(a) Why are alkali and alkaline earth metals among the most reactive elements of the periodic table?

  • Alkali metals (Group 1) and alkaline earth metals (Group 2) have low ionization energies and readily lose electrons to form cations. Their reactivity increases down the group due to the decreasing ionization energy.

(b) Why does lime water turn milky with CO₂ but becomes clear with excess CO₂?

  • When CO₂ is bubbled through lime water (Ca(OH)₂), it reacts to form calcium carbonate (CaCO₃), which is insoluble and causes the milky appearance:
    [ Ca(OH)₂ + CO₂ &→ CaCO₃ + H₂O ]
    On passing excess CO₂, calcium carbonate reacts with water and CO₂ to form soluble calcium bicarbonate (Ca(HCO₃)₂), causing the solution to clear:
    [ CaCO₃ + CO₂ + H₂O → Ca(HCO₃)₂ ]

(c) How is gypsum converted into plaster of Paris?

  • Gypsum (CaSO₄·2H₂O) is heated at around 150°C to 180°C, losing water and converting into plaster of Paris (CaSO₄·½H₂O):
    [ CaSO₄·2H₂O →{\text{150°C-180°C}} CaSO₄·½H₂O + 1½H₂O ]

(d) Why is 2% gypsum added in cement?

  • Gypsum is added to cement to slow down the setting process by controlling the hydration reaction of the cement. It prevents the cement from setting too quickly during the mixing and application phase.

(e) Why is lime added to acidic soil?

  • Lime (CaO or Ca(OH)₂) neutralizes the excess acidity in soil by reacting with the acids, thus increasing the pH and improving soil fertility.

(f) How are lime and sand used to make glass?

  • Lime (CaO) and sand (SiO₂) are melted together at high temperatures to form calcium silicate (CaSiO₃), a component of glass. Lime acts as a stabilizer, preventing the glass from being too soluble in water.

(g) How is lime mortar prepared?

  • Lime mortar is prepared by mixing slaked lime (Ca(OH)₂) with sand and water. When exposed to air, the lime reacts with carbon dioxide (CO₂) to form calcium carbonate, hardening the mortar:
    [ Ca(OH)₂ + CO₂ → CaCO₃ + H₂O ]

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