Chapter 9

Explain the contributions of Johann Döbereiner and John Newlands to the organization of elements according to their properties.

Who is credited with arranging the periodic table? How are the elements arranged in the modern periodic table?

What is an electron configuration? Give an example.

What is Coulomb's law? Explain how the potential energy of two charged particles depends on the distance between the charged particles and on the magnitude and sign of their charges.

What is shielding? In an atom, which electrons tend to do the most shielding (core electrons or valence electrons)?

What is penetration? How does the penetration of an orbital into the region occupied by core electrons affect the energy of an electron in that orbital?

Why are the sublevels within a principal level split into different energies for multielectron atoms but not for the hydrogen atom?

Why is electron spin important when writing electron configurations? Explain in terms of the Pauli exclusion principle.

What are degenerate orbitals? According to Hund's rule, how are degenerate orbitals occupied?

List all orbitals from \(1 s\) through \(5 s\) according to increasing energy for multielectron atoms.

What are valence electrons? Why are they important?

Copy this blank periodic table onto a sheet of paper and label each of the blocks within the table: \(s\) block, \(p\) block, \(d\) block, and \(f\) block.

Explain why the \(s\) block in the periodic table has only two columns while the \(p\) block has six.

Why do the rows in the periodic table get progressively longer as you move down the table? For example, the first row contains 2 elements, the second and third rows each contain 8 elements, and the fourth and fifth rows each contain 18 elements. Explain.

Describe the relationship between a main-group element's lettered group number (the number of the element's column) and its valence electrons.

Describe the relationship between an element's row number in the periodic table and the highest principal quantum number in the element's electron configuration. How does this relationship differ for main-group elements, transition elements, and inner transition elements?

Which of the transition elements in the first transition series have anomalous electron configurations?

Describe how to write the electron configuration for an element based on its position in the periodic table.

Describe the relationship between the properties of an element and the number of valence electrons that it contains.

List the number of valence electrons for each family, and explain the relationship between the number of valence electrons and the resulting chemistry of the elements in the family. a. alkali metals b. alkaline earth metals c. halogens d. oxygen family

Define atomic radius. For main-group elements, describe the observed trends in atomic radius as you move: a. across a period in the periodic table b. down a column in the periodic table

Use the concepts of effective nuclear charge, shielding, and \(n\) value of the valence orbital to explain the trend in atomic radius as you move across a period in the periodic table.

For transition elements, describe and explain the observed trends in atomic radius as you move: a. across a period in the periodic table b. down a column in the periodic table

How is the electron configuration of an anion different from that of the corresponding neutral atom? How is the electron configuration of a cation different?

Describe how to write an electron configuration for a transition metal cation. Is the order of electron removal upon ionization \(\operatorname{sim}-\) ply the reverse of electron addition upon filling? Why or why not?

Describe the relationship between a. the radius of a cation and that of the atom from which forms b. the radius of an anion and that of the atom from which it forms

What is ionization energy? What is the difference between first ionization energy and second ionization energy?

What is the general trend in first ionization energy as you move down a column in the periodic table? As you move across a row?

What is electron affinity? What are the observed periodic trends in electron affinity?

Write the full electron configuration for each element. a. Si b. O c. K d. Ne

Write the full electron configuration for each element. a. C b. P c. Ar d. Na

Write the full orbital diagram for each element. a. \(\mathrm{N}\) b. F c. Mg d. Al

Write the full orbital diagram for each element. a. S b. Ca c. Ne d. He

Use the periodic table to write an electron configuration for each element. Represent core electrons with the symbol of the previous noble gas in brackets. a. \(P\) b. Ge c. Zr d. I

Use the periodic table to determine the element corresponding to each electron configuration. a. \([\mathrm{Ar}] 4 s^{2} 3 d^{10} 4 p^{6}\) b. \([\mathrm{Ar}] 4 s^{2} 3 d^{2}\) c. \([\mathrm{Kr}] 5 s^{2} 4 d^{10} 5 p^{2}\) d. \([\mathrm{Kr}] 5 s^{2}\)

Use the periodic table to determine each quantity. a. the number of \(2 s\) electrons in Li b. the number of \(3 d\) electrons in Cu c. the number of \(4 p\) electrons in Br d. the number of \(4 d\) electrons in \(\mathrm{Zr}\)

Use the periodic table to determine each quantity. a. the number of \(3 s\) electrons in \(\mathrm{Mg}\) b. the number of \(3 d\) electrons in Cr c. the number of \(4 d\) electrons in Y d. the number of \(6 p\) electrons in \(\mathrm{Pb}\)

Name an element in the fourth period (row) of the periodic table with the following: a. five valence electrons b. four \(4 p\) electrons c. three \(3 d\) electrons d. full \(s\) and \(p\) sublevels

Name an element in the third period (row) of the periodic table with the following: a. three valence electrons b. four \(3 p\) electrons c. \(\operatorname{six} 3 p\) electrons d. two \(3 s\) electrons and no \(3 p\) electrons

Determine the number of valence electrons in an atom of each element. a. Ba b. Cs C. \(\mathrm{Ni}\) d. S

Which outer electron configuration would you expect to belong to a reactive metal? To a reactive nonmetal? a. \(n s^{2}\) b. \(n s^{2} n p^{6}\) c. \(n s^{2} n p^{5}\) d. \(n s^{2} n p^{2}\)

Which outer electron configurations would you expect to belong to a noble gas? To a metalloid? a. \(n s^{2}\) b. \(n s^{2} n p^{6}\) c. \(n s^{2} n p^{5}\) d. \(n s^{2} n p^{2}\)

According to Coulomb's law, which pair of charged particles has the lowest potential energy? a. a particle with a \(1-\) charge separated by \(150 \mathrm{pm}\) from a particle with a \(2+\) charge b. a particle with a \(1+\) charge separated by \(150 \mathrm{pm}\) from a particle with a \(1+\) charge c. a particle with a \(1-\) charge separated by \(100 \mathrm{pm}\) from a particle with a \(3+\) charge

According to Coulomb's law, rank the interactions between charged particles from lowest potential energy to highest potential energy. a. a \(1+\) charge and a \(1-\) charge separated by \(100 \mathrm{pm}\) b. a \(2+\) charge and a \(1-\) charge separated by \(100 \mathrm{pm}\) c. a \(1+\) charge and a \(1+\) charge separated by \(100 \mathrm{pm}\) d. a \(1+\) charge and a \(1-\) charge separated by \(200 \mathrm{pm}\)

Which experience a greater effective nuclear charge: the valence electrons in beryllium or the valence electrons in nitrogen? Why?

Arrange the atoms according to decreasing effective nuclear charge experienced by their valence electrons: \(\mathrm{S}, \mathrm{Mg}, \mathrm{Al}, \mathrm{Si} .\)

Choose the larger atom from each pair, if possible. a. Sn or \(\mathrm{Si}\) b. Br or Ga c. Sn or Bi d. Se or \(\mathrm{Sn}\)

Arrange these elements in order of increasing atomic radius: \(\mathrm{Ca}, \mathrm{Rb}, \mathrm{S}, \mathrm{Si}, \mathrm{Ge}, \mathrm{F}\)

Arrange these elements in order of decreasing atomic radius: \(\mathrm{Cs}, \mathrm{Sb}, \mathrm{S}, \mathrm{Pb}, \mathrm{Se}\)

Write the electron configuration for each ion. a. \(\mathrm{O}^{2-}\) b. \(\mathrm{Br}^{-}\) c. \(\mathrm{Sr}^{2+}\) d. \(\mathrm{Co}^{3+}\) e. \(\mathrm{Cu}^{2+}\)