Free s-Block and p-Block Elements MCQs with Answers

1. Across a period from left to right, the atomic radius of the elements generally

• A. decreases, because nuclear charge increases while electrons enter the same shell
• B. increases, because more electrons are added
• C. stays the same
• D. increases, because the nuclear charge decreases

Explanation: Each successive element adds one proton and one electron, but the electron enters the same outermost shell, so shielding hardly changes while the effective nuclear charge rises and the shell is pulled inwards. Down a group the opposite happens, since a whole new shell is added and radius increases. Atomic radius is the single trend from which ionisation energy and electronegativity mostly follow.

2. The first ionisation energy is the energy required to

• A. add one electron to a gaseous atom
• B. remove one mole of electrons from one mole of gaseous atoms, forming gaseous unipositive ions
• C. break one mole of covalent bonds
• D. convert one mole of a solid into a gas

Explanation: The definition specifies the gaseous state so that no lattice or intermolecular forces interfere, and it is always endothermic because an electron must be pulled away from an attracting nucleus. Adding an electron to a gaseous atom is electron affinity, the reverse process. Second ionisation energy is always larger than the first, since the electron is being removed from a positively charged ion.

3. Which element has the highest electronegativity?

• A. Oxygen
• B. Chlorine
• C. Fluorine
• D. Nitrogen

Explanation: Fluorine is the most electronegative element on the Pauling scale at 4.0, because it is small and has a high effective nuclear charge, so it attracts bonding electrons more strongly than any other atom. Oxygen at 3.5 is second and chlorine and nitrogen follow at 3.0. Electronegativity increases across a period and decreases down a group, so the maximum lies at the top right, excluding the noble gases.

4. Down group I from lithium to caesium, the reactivity of the metals

• A. decreases, because atomic radius increases
• B. stays constant
• C. decreases, because ionisation energy falls
• D. increases, because the outer electron is further from the nucleus and more easily lost

Explanation: Each element down the group has an extra shell, so the single outer electron is further from the nucleus and better shielded, ionisation energy falls and the metal loses that electron more readily. This is why potassium ignites the hydrogen it releases from water while lithium merely fizzes. In the halogens the same reasoning gives the opposite trend, since there reactivity depends on gaining an electron.

5. The elements of group I are called the alkali metals because they

• A. are found free in nature
• B. react with water to form hydroxides that are strongly basic
• C. are all radioactive
• D. form acidic oxides

Explanation: Sodium and its neighbours react with water to give hydrogen and a hydroxide solution that is strongly alkaline, which is the origin of the family name. Because they are so reactive they are never found uncombined in nature and are stored under oil. Their oxides are basic, not acidic, which is typical of metals.

6. When sodium reacts with excess oxygen on burning, the main product is

• A. sodium oxide, Na2O
• B. sodium superoxide, NaO2
• C. sodium peroxide, Na2O2
• D. sodium hydroxide, NaOH

Explanation: Lithium forms the simple oxide, sodium forms mainly the peroxide and potassium, rubidium and caesium form superoxides, the larger cations being better able to stabilise the larger anions. Sodium peroxide reacts with water to give hydrogen peroxide and alkali. Sodium hydroxide is the product of reaction with water rather than with oxygen.

7. Group II elements are called alkaline earth metals and, compared with group I, they are

• A. less reactive, with higher melting points and harder solids
• B. more reactive in every case
• C. non metallic in character
• D. monovalent in their compounds

Explanation: Group II atoms are smaller and must lose two electrons rather than one, so more energy is needed and they are less reactive than their group I neighbours. Having two delocalised electrons per atom also gives stronger metallic bonding, hence higher melting points and greater hardness. They are divalent, forming ions with a charge of plus 2.

8. The thermal stability of the group II carbonates from magnesium to barium

• A. decreases down the group
• B. is the same for all
• C. increases down the group
• D. first increases then decreases

Explanation: A small, highly charged cation polarises the carbonate ion strongly and weakens the carbon to oxygen bonds, so magnesium carbonate decomposes at a relatively low temperature while barium carbonate needs a much higher one. Polarising power falls as the cation gets larger down the group, so stability rises. The nitrates and hydroxides follow the same pattern for the same reason.

9. Which of the following group II compounds is used in medicine as an antacid?

• A. Barium sulphate
• B. Magnesium hydroxide
• C. Calcium carbide
• D. Beryllium chloride

Explanation: Magnesium hydroxide, sold as milk of magnesia, is a weak insoluble base that neutralises excess stomach acid without being absorbed in quantity. Barium sulphate is also used medically but as a contrast medium for X rays, and it is safe only because it is so insoluble that the toxic barium ions are not released. Beryllium compounds are highly toxic.

10. The elements of group IV show a gradation in character from

• A. metallic carbon to non metallic lead
• B. non metallic carbon through metalloid silicon and germanium to metallic tin and lead
• C. entirely metallic throughout
• D. entirely non metallic throughout

Explanation: Carbon is a non metal, silicon and germanium are metalloids used as semiconductors, and tin and lead are true metals, so group IV shows the clearest metallic gradation of any group. The reason is that ionisation energy falls down the group as the atoms get larger. This trend of increasing metallic character down a group applies across the p block.