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).
