chem Ch 3

1. According to Coulomb's law, if the separation between two particles of the same charge is doubled, the potential energy of the two particles ________________.

. is one-half as high as it was before the separation

A main-group element has an outer electron configuration of ns2np4. What charge is likely for an ion of this element?

2-

On the basis of periodic trends, choose the larger atom in each pair (if possible). Explain your choices. C or F C or Ge N or Al Al or Ge

C atoms are larger than F atoms because as you trace the path between C and F on the periodic table, you move to the right within the same period. As you move to the right across a period, the effective nuclear charge experienced by the outermost electrons increases, resulting in a smaller radius. Ge atoms are larger than C atoms because as you trace the path between C and Ge on the periodic table, you move down a column. Atomic size increases as you move down a column because the outermost electrons occupy orbitals with a higher principal quantum number that are therefore larger, resulting in a larger atom. Al atoms are larger than N atoms because as you trace the path between N and Al on the periodic table, you move down a column (atomic size increases) and then to the left across a period (atomic size increases). These effects add together for an overall increase. Based on periodic trends alone, you cannot tell which atom is larger because as you trace the path between Al and Ge you move to the right across a period (atomic size decreases) and then down a column (atomic size increases). These effects tend to counter each other, and it is not easy to tell which will predominate.

. In the previous sections, we have seen how the number of electrons and the number of protons affect the size of an atom or ion. However, we have not considered how the number of neutrons affects the size of an atom. Would you expect isotopes, for example, C-12 and C-13, to have different atomic radii?

C-12 and C-13 are the same size.

. Which statement is true about electron shielding of nuclear charge?

Core electrons efficiently shield outermost electrons from nuclear charge.

5. For which element is the gaining of an electron most exothermic?

F

Write the electron configuration for Ge. Identify the valence electrons and the core electrons.

Ge 1s22s22p63s23p64s23d104p2

Classify the bond formed between each pair of atoms as covalent, polar covalent, or ionic. Sr and F N and Cl N and O

Ionic, covalent, polar covalent

1. Which compound is most likely to contain ionic bonds? a. CH4 b. N2O c. MgF2

MgF2

On the basis of periodic trends, determine which element in each pair has the higher first ionization energy (if possible). Al or S As or Sb N or Si O or Cl

S has a higher ionization energy than Al because as you trace the path between Al and S on the periodic table, you move to the right within the same row. Ionization energy increases as you go to the right due to increasing effective nuclear charge. As has a higher ionization energy than Sb because as you trace the path between As and Sb on the periodic table, you move down a column. Ionization energy decreases as you go down a column as a result of the increasing size of orbitals with increasing n. N has a higher ionization energy than Si because as you trace the path between N and Si on the periodic table, you move down a column (ionization energy decreases) and then to the left across a row (ionization energy decreases). These effects sum together for an overall decrease. Based on periodic trends alone, it is impossible to tell which has a higher ionization energy because, as you trace the path between O and Cl, you go to the right across a row (ionization energy increases) and then down a column (ionization energy decreases). These effects tend to counter each other, and it is not obvious which will dominate.

The ionization energies of an unknown third period element are shown here. Identify the element. IE1 = 786 kJ/mol; IE2 = 1580 kJ/mol; IE3 = 3230 kJ/mol; IE4 = 4360 kJ/mol; IE5 = 16,100 kJ/mol;

Si

On the basis of periodic trends, select the larger atom: Sn or I.

Sn

On the basis of periodic trends, choose the more metallic element from each pair (if possible). Sn or Te P or Sb Ge or In S or Br

Sn is more metallic than Te because as you trace the path between Sn and Te on the periodic table, you move to the right within the same period. Metallic character decreases as we move to the right. Sb is more metallic than P because as you trace the path between P and Sb on the periodic table, you move down a column. Metallic character increases as we move down a column. In is more metallic than Ge because as you trace the path between Ge and In on the periodic table, you move down a column (metallic character increases) and then to the left across a period (metallic character increases). These effects add together for an overall increase. Based on periodic trends alone, we cannot tell which is more metallic because as you trace the path between S and Br, you move to the right across a period (metallic character decreases) and then down a column (metallic character increases). These effects tend to counter each other, and it is not easy to tell which will predominate.

According to Coulomb's law, what happens to the potential energy of two oppositely charged particles as they get closer together?

Their potential energy decreases.

Write electron configurations for each element. Mg P Br Al

[Ne]3s^2 [Ne]3s^23p^3 [Ar]3d^104s^24p^5 [Ne]3s^23p^1

1. Based on what you just learned about ionization energies, explain why valence electrons are more important than core electrons in determining the reactivity and bonding in atoms.

a. Since bonding involves the transfer or sharing of electrons, valence electrons are most important because they are held most loosely.

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

a. Sodium has only one valence electron that is easily removed; removing a second electron would be more difficult because it would have to be a core electron.

1. What are the four quantum numbers for each of the two electrons in a 4s orbital?

a. n = 4, l = 0, ml = 0, ms = +½; n = 4, l = 0, ml = 0, ms = −½

Which statement is true about effective nuclear charge?

b. Effective nuclear charge increases as we move to the right across a row in the periodic table.

4. Identify the correct trends in metallic character.

b. Metallic character decreases as we move to the right across a row in the periodic table and increases as we move down a column.

. Arrange these elements in order of increasing first ionization energy: Cl, Sn, Si.

b. Sn < Si < Cl

0. Which species is diamagnetic?

b. Zn

Which statement is true? a. An orbital that penetrates into the region occupied by core electrons is more shielded from nuclear charge than an orbital that does not penetrate and therefore has a higher energy. b. An orbital that penetrates into the region occupied by core electrons is less shielded from nuclear charge than an orbital that does not penetrate and therefore has a higher energy. c. An orbital that penetrates into the region occupied by core electrons is less shielded from nuclear charge than an orbital that does not penetrate and therefore has a lower energy. d. An orbital that penetrates into the region occupied by core electrons is more shielded from nuclear charge than an orbital that does not penetrate and therefore has a lower energy.

c. An orbital that penetrates into the region occupied by core electrons is less shielded from nuclear charge than an orbital that does not penetrate and therefore has a lower energy.

Arrange these atoms and ions in order of increasing radius: Cs+, Ba2+, I-.

c. Ba2+ < Cs+ < I-

What is the electron configuration for Fe2+?

c. [Ar]4s0 3d6

Which electron in S is most shielded from nuclear charge? 1s, 2p, or 3p

electron in a 3p orbital (valence electrons more shielded)