Q: How is the effective nuclear charge ($$Z_{\text{eff}}) $$calculated for a given electron?

$$Z_{\text{eff}} = Z - S$$, where $$Z is $$the atomic number and $$S is $$the number of core electrons

Question 1

Which of the following best describes the main concept that determines the properties of atoms?

Accepted Answer

The arrangement of electrons around the nucleus

Question 2

Effective nuclear charge ($$Z_{	ext{eff}}) is $$best defined as:

Accepted Answer

The net positive charge experienced by an electron in a multi-electron atom

Question 3

How is the effective nuclear charge ($$Z_{	ext{eff}}) $$calculated for a given electron?

Accepted Answer

$$Z_{	ext{eff}} = Z - S$$, where $$Z is $$the atomic number and $$S is $$the number of core electrons

Question 4

Which element in the following set is expected to have the highest effective nuclear charge ($$Z_{	ext{eff}}) $$for its valence electrons: Na, Mg, Al, Si?

Accepted Answer

Si

Question 5

What is the general trend for atomic radius as you move across a period from left to right in the periodic table, and why does this occur?

Accepted Answer

Atomic radius decreases due to increased effective nuclear charge

Question 6

Compare the size of a sodium atom (Na) to a sodium ion (Na$$^+). $$Which is larger and why?

Accepted Answer

Na is larger because it has more electrons and less effective nuclear charge per electron

Question 7

Which of the following sets contains ions that are isoelectronic?

Accepted Answer

O$$^{2-}$$, F$$^-$$, Na$$^+$$, Mg$$^{2+}$$

Question 8

Which of the following best defines ionization energy?

Accepted Answer

The energy required to remove an electron from a gaseous atom

Question 9

What is the trend for first ionization energy as you move down a group in the periodic table?

Accepted Answer

It decreases due to increased atomic radius and electron shielding

Question 10

Which of the following is the correct electron configuration for the chloride ion (Cl$$^-$$)?

Accepted Answer

$$1s^2$$\ $$2s^2$$\ $$2p^6$$\ $$3s^2$$\ $$3p^6$$

Question 11

Which statement best describes the difference between electron affinity (EA) and ionization energy (IE)?

Accepted Answer

EA is the energy released when an atom gains an electron, while IE is the energy required to remove an electron

Question 12

Which element is expected to have the most negative electron affinity (i.e., releases the most energy when gaining an electron)?

Accepted Answer

Fluorine (F)

Question 13

Which of the following best describes the trend in metallic character across a period from left to right?

Accepted Answer

Metallic character decreases

Question 14

Which of the following is NOT a typical property of metals?

Accepted Answer

High ionization energy

Question 15

Which of the following is a property commonly associated with nonmetals?

Accepted Answer

Brittle in solid form

Question 16

Which property is unique to Group 1A (alkali) metals?

Accepted Answer

They form $$+1$$ ions and react vigorously with water

Question 17

Which property is unique to Group 2A (alkaline earth) metals compared to Group 1A metals?

Accepted Answer

They form $$+2$$ ions and are less reactive with water

Question 18

Why are the trends for the selected elements in the nonmetallic section of the periodic table (such as the halogens and noble gases) considered special?

Accepted Answer

They have completely filled s and p orbitals, leading to unique chemical inertness or high reactivity

Question 19

A patient is exposed to a toxic dose of sodium metal. Which property of sodium explains its dangerous reaction with water in the body?

Accepted Answer

Sodium forms $$+1$$ ions and reacts vigorously with water to produce hydrogen gas and heat

Question 20

Which of the following best explains why fluorine is more reactive than chlorine?

Accepted Answer

Fluorine has a higher effective nuclear charge and smaller atomic radius, increasing its tendency to gain electrons

Chapter 7 Study Guide: Periodic Properties of the Elements

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