The atomic radius of sodium is 186 pm and of chlorine is 100 pm. The ionic radius for Na + is 102 pm and for Cl - is 181 pm. In going from Na to Cl in Period 3, why does the atomic radius decrease while the ionic radius increases? a. The inner electrons in the sodium cation shield its valence electrons more effectively than the inner electrons in the chloride anion do. b. The inner electrons shield the valence electrons more effectively in the chlorine atom than in the chloride anion. c. The outermost electrons in chloride experience a smaller effective nuclear charge than those in the sodium cation do. d. The outermost electrons in chloride experience a larger effective nuclear charge than those in the sodium cation do. e. Monatomic ions are bigger than the atoms from which they are formed.

Question

Chemistry

Marques Welch

8 June, 23:11

The atomic radius of sodium is 186 pm and of chlorine is 100 pm. The ionic radius for Na + is 102 pm and for Cl - is 181 pm. In going from Na to Cl in Period 3, why does the atomic radius decrease while the ionic radius increases? a. The inner electrons in the sodium cation shield its valence electrons more effectively than the inner electrons in the chloride anion do. b. The inner electrons shield the valence electrons more effectively in the chlorine atom than in the chloride anion. c. The outermost electrons in chloride experience a smaller effective nuclear charge than those in the sodium cation do. d. The outermost electrons in chloride experience a larger effective nuclear charge than those in the sodium cation do. e. Monatomic ions are bigger than the atoms from which they are formed.

+1

Answers (1)

Know the Answer?

Andy Guerra · Answer 1

The outermost electrons in chloride experience a smaller effective nuclear charge than those in the sodium cation do

Explanation:

Effective nuclear charge refers to the attractive positive charge of nuclear protons acting on valence electrons. The effective nuclear charge is always less than the total number of protons present in a nucleus due to shielding effect. Effective nuclear charge is behind all other periodic table tendencies. The effective nuclear charge is the net positive charge experienced by valence electrons. It can be approximated by the equation: Zeff = Z - S, where Z is the atomic number and S is the number of shielding.

The atomic number increases moving left to right across a period and so does the effective nuclear charge likewise. Therefore, moving left to right across a period the nucleus has a greater pull on the outer electrons and the atomic radii decreases. This explains the fact that the formation of a positive ion leads to an increase in effective nuclear charge since the number of electrons reduces and the nucleus pulls the outermost shell closer to the nucleus hence the cation decreases in size. Similarly, the formation of an anion leads to an increase in ionic radius owing to the fact that effective nuclear charge decreases because of the addition of an electron. Hence the radius of an anion is greater than the radius of a cation.