Question 1

Which of the following best explains why the first ionization energy generally increases as you move from left to right across a period in the periodic table?

Accepted Answer

The effective nuclear charge ($$Z_{\mathrm{eff}}) $$increases, pulling electrons closer to the nucleus.

Question 2

Which element in the following pair has the higher first ionization energy: sulfur (S) or aluminum (Al)?

Accepted Answer

Sulfur (S), because it is farther to the right in the same period.

Question 3

Which of the following equations correctly represents the second ionization energy for magnesium?

Accepted Answer

$$\mathrm{Mg}^+(g) + IE_2$ \rightarrow \mathrm{Mg}^{2+}(g) + 1e^-$$

Question 4

Why is the first ionization energy of boron (B) lower than that of beryllium (Be), even though boron is to the right of beryllium in the same period?

Accepted Answer

Boron’s valence electron is in a 2p orbital, which is less shielded and easier to remove than beryllium’s 2s electron.

Question 5

A second-row element shows a large jump between its third and fourth ionization energies. Which element is it most likely to be?

Accepted Answer

Aluminum (Al), because after removing three electrons, the next electron comes from a core shell.

Question 6

Which statement best explains why sodium commonly forms a $$1^+$$ ion and not a $$2^+$$ ion?

Accepted Answer

Sodium has only one valence electron that is easily removed; removing a second electron would be a core electron.

Question 7

Which of the following trends is correct for electron affinity (EA) among main-group elements?

Accepted Answer

EA becomes more negative (more exothermic) as you move to the right across a period.

Question 8

Which group of elements generally has the most negative electron affinity values?

Accepted Answer

Halogens (Group 7A)

Question 9

Which of the following statements about metallic character is true?

Accepted Answer

Metallic character increases as you move down a group.

Question 10

Arrange the following elements in order of increasing metallic character: Si, Cl, Na, Rb.

Accepted Answer

Cl \u003c Si \u003c Na \u003c Rb

Question 11

Which of the following best describes the relationship between atomic radius and ionization energy down a group?

Accepted Answer

As atomic radius increases, ionization energy decreases.

Question 12

Which of the following elements is more metallic: tin (Sn) or tellurium (Te)?

Accepted Answer

Tin (Sn), because metallic character decreases as you move to the right across a period.

Question 13

Which of the following is a property of metals but NOT of nonmetals?

Accepted Answer

Conduct electricity

Question 14

Why do alkali earth metals and noble gases generally have low (or positive) electron affinities?

Accepted Answer

Their added electron goes into a higher sublevel or energy level, making the process less exothermic.

Question 15

Which of the following best explains why valence electrons are more important than core electrons in determining the reactivity and bonding in atoms?

Accepted Answer

Valence electrons are held most loosely and are involved in bonding.

Question 16

Which of the following statements about successive ionization energies is correct?

Accepted Answer

There is a large increase in ionization energy when a core electron is removed after all valence electrons.

Question 17

Which of the following elements would you expect to have the lowest first ionization energy?

Accepted Answer

Sodium (Na)

Question 18

Which of the following best describes the general trend in metallic character across a period?

Accepted Answer

Metallic character decreases from left to right.

Question 19

Which of the following is the correct equation for the electron affinity of chlorine?

Accepted Answer

$$\mathrm{Cl}(g) + 1e^- \rightarrow \mathrm{Cl}^-(g) + EA$$

Question 20

Which of the following elements would you expect to have the most exothermic electron affinity?

Accepted Answer

Chlorine (Cl)

lectuer 13

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