Study the table 1 and 2 related to electronegativity values along the period 2 and 3 Group 1 and 17 and answer the questions that follow based on these.
Electronegativity Values (on Pauling scale) Across the Periods

Electronegativity Values (on Pauling scale) Down a Family

(a) How does electronegativity varies along a period and why?
(b) Why does K and Rb have same electronegativity although Rb is bigger than K?
(c) Why does electronegativity decrease down the group?
(d) Why is CsF most ionic?
(e) Why is H- ion bigger than F^- ion?
(f) Why is Lil covalent?
(g) Why is CI₂ O₇ more acidic than Cl₂O₅?

Study the table of electron gain enthalpies of some main group elements and answer the questions that follow based on table and related concepts.

(a) Why do group 1 elements have low electron gain enthalpies?
(b) Why is electron gain enthalpy of sulphur more than oxygen?
(c) Why do noble gases have +ve electron gain enthalpies?
(d) Arrange P, S, CI, F is increasing order of electron gain enthalpeis.

Observe the group shown in figure and answer the questions that follow based on the graph and related studied concepts.

(a) Which group elements have lowest ionisation enthalpy and why?
(b) Which group element have highest ionisation enthalpy and why?
(c) Why is ionisation enthalpy of Be more than B?
(d) Why does 'N' have higher ionisation enthalpy than O?

The atomic and ionic radii decrease with increase in atomic number along a period from left to right. Atomic size and ionic size increases down the group.
Ionisation enthalpy decreases down the group and increases along a period from left to right. It also depends upon shielding effect as well as stability of electronic configuration.
Electronegativity decreases down the group but increases along the period.
(a) Arrange the elements of second period in increasing order of first ionisation enthalpy.
(b) Arrange the elements of group 13 in increasing order of atomic size.
(c) Select the' amphoteric oxides among the following:
NO, B₂O₃, Cr ₂O₃, BeO, ZnO, CO₂ , Al₂ O₃
(d) For an element I.E₁ = 738 kJ mol^-1, IE₂ = 1450 kJ mol^-1, IE₃ = 7700 kJ mol^-1, IE₄ = 1.1 x 10 ⁴ kJ mol^-1. Name the main group to which element belong. Why?

Modern periodic table arranges the elements in the increasing order of atomic number. It has 18 groups and 7 periods. Atomic numbers are consecutive in a period and increases in group in a pattern. Elements are divided into four blocks, s-block,p-block, d-block andf-block based on their electronic configuration. 78% of elements are metals, about 20 elements are non-metals and few elements like B, Si, Ge, As are metalloids. Metallic character increases down the group but decreases along the period from left to right. The physical and chemical properties vary periodically with their atomic numbers.
Periodic trends are observed in atomic size, ionisation enthalpies, electron gain enthalpies, electronegativity and valence. Oxides of metals are basic, some are amphoteric. Non-metals form acidic oxides, some form neutral oxides. s-block elements are soft, highly reactive, do not show variable oxidation states. p-block elements are metals, non-metals as well as metalloids, show variable oxidation states, exist as solids, liquids and gases. d-block elements are metals, form coloured ions, show variable oxidation states, have high melting and boiling points. Lanthanoids and actinoids are f-block elements, form coloured ions. All actinoids are radioactive.
(a) Name the elements which belong to d-block but are not transition metals.
(b) What are representative elements?
(c) What is difference between oxidation states of p-block and d-block elements?
(d) Which group elements are most electropositive and why?
(e) What happens to reactivity down the group 17?
(f) What type of compounds are formed by element belonging to group 14 and third period? Identify the element.
(g) What is formula of compound formed between AI and S