Chem 4.4
Based on your knowledge of trends in the periodic table, place the following atoms in order of increasing ionization energy: Ar, Sr, P, Mg, and Ba.
Ba < Sr < Mg< P < Ar
Why are ionization energies so much higher when core electrons are involved than they are when only valence elec- trons are involved?
Core electrons feel the full attraction of the nucleus. Valence electrons are held much less tightly because of the shield effect of the core electrons screening off the attraction of the positive nucleus from the valence electrons.
Using only the periodic table as a reference (without electronegativity data), predict the relative electronegativites of these elements and put them in order from least to greatest: Ni, Ta, Se, F, Cs, Cl.
Cs < Ta < Ni < Se < Cl < F
Write the condensed electron configurations for Cu²⁺, As⁵⁺, Ag⁺, and Au³⁺.
Cu²⁺: [Ar]3d⁹ As⁵⁺: [Ar]3d^10 Ag⁺: [Kr]4d^10 Au³⁺: [Xe]4f^14
Aluminum and scandium both ionize to +3, even though scandium is in Group 3 and aluminum is in Group 13. Explain why this is.
Each has 3 valence electrons. Aluminum is third from the left in Period 3, scandium is third from the left in Period 4.
Distinguish between electron affinity and electronegativity.
Electron affinity is the energy released when an electron is added to an atom. Electronegativity is the relative strength of attrac- tion of an atom for shared electron pairs in molecular bonds.
Distinguish between ionization energy and electron affinity.
Ionization energy is the energy required to remove an electron from an atom. Electron affinity is the energy released when an electron is added to an atom.
Define ionization energy and describe the trends for ionization energy in the periodic table across periods and down groups.
Ionization energy is the energy required to remove an electron from an atom. This energy increases from left to right across the table, and decreases down the groups.
Why do the chalcogens form ions with a charge of -2?
Main group elements ionize so as to end up with either an empty or full valence shell. Chalcogens need two more electrons to have a full valence shell, so they ionize by gaining two electrons, acquiring a change of -2 in the process.
Based on your knowledge of trends in the periodic table, place the following atoms in order of increasing electron affinity: Br, Rb, and S.
Rb < S < Br
Which of the following is likely to have the greatest difference between the third and fourth ionization energies: Cl, Sc, Na, C?
Scandium (Sc), due to the fact that it has 3 valence electrons, so loses these easily. A fourth ionization would require losing a core electron, which requires a much higher energy to accomplish.
Of the following cations, which is least likely to form: Ca³₊, Mg²⁺, K⁺? Explain your response.
The Ca³⁺ is least likely. Ca normally ionizes to Ca³⁺. To ionize to Ca³⁺ a core electron would have to be removed from the atom, which does not normally occur.
Develop an explanation for why the electron affinity values for the chalcogens are each significantly lower than those of their halogen neighbors.
The further to the right an element is in the table, the less pronounced the shield effect is. The shield effect reduces the attraction of a nucleus for the atom's electrons, but also affects the attraction of the nucleus for other atoms' electrons. Halogens are to the right of the chalcogens, so the shield effect in the halogens is less pronounced and they thus their nuclei attract neigh- boring atoms' electrons more strongly. This means they have higher electronegativities.
Referring again to Table 4.1, explain the large increase in ionization energy that occurs in the yellow shaded region of that table.
The high ionization energies are for core electrons, which are bound much more tightly to the nucleus than the valence elec- trons are.
Explain what the electronegativity scale is used for and how it arose.
The scale was invented by Linus Pauling in 1932 while he was researching chemical bonding. The scale describes the relative strength of the attraction of atoms of different elements for shared electrons in molecular bonds. Differences in electronegativity account for the polarization that occurs in molecules due to the different strength of attraction in the molecule of different atoms for the shared electrons in the chemical bonds.
Write a description accounting for the trends in ionization energy in terms of the shield effect and other factors.
The strong upward trend moving across each period is accounted for by the decreasing efficacy of the shield effect that occurs moving from left to right in the period. Moving from left to right, the strength of the shielding provided by the atom's core electrons remains the same, but the positive charge of the nucleus increases with each additional proton. This means the shield effect becomes less effective and the attraction between the nucleus and the valence electrons increases from left to right. The higher the nuclear attraction is, the stronger the nucleus attracts the outermost electrons and the greater the energy required to remove one of them. The trend down a group is accounted for by the fact that the elements in each new period have their valence electrons in a shell with a higher principle quantum number. These electrons are thus farther from the nucleus, have a less negative energy, and are easier to remove.
Why don't the noble gases have electronegativity values listed in Figure 4.15?
These elements have full valence shells and are unreactive. They do not normally ionize.
Referring to the periodic table, describe the chemical properties of potassium (K), sulfur (S), xenon (Xe), iodine, (I), and manganese (Mn).
potassium: has 1 valence electron, ionizes to +1, has a very low electronegativity sulfur: has 6 valence electrons, ionizes to -2, has a very high electronegativity xenon: is a noble gas, so has a full valence shell and is unreactiveiodine: has 7 valence electrons, ionizes to -1, has a very high electronegativity manganese: has 2 valence electrons, ionizes to +2, has average electronegativity
