Periodic Trends: Size Trends in Atomic Radii Imagine transforming lithium to beryllium in two steps: 1. Add an electron to Li. 2. Then, add a proton (plus neutrons, but they don t count) to the nucleus. As we go across the rows of the periodic table, the electron-electron repulsion from adding an electron to a same-shell orbital (2s, 2p, 3s, 3p, etc.) is a decrease in effective charge that is always less than the increase in nuclear charge. 1
Ionic Radii Adding or removing electrons from an atom will always change its size: These effects become quite large when adding an electron or removing it empties a valence shell (Na Na + ) or starts a new one (Xe Xe.) Isoelectronic Series Once we add or remove electrons to form ions, we can compare multiple species (anions, neutrals, cations) that have different nuclear charges but the same electronic configuration. Here are the electronic configurations of oxygen, fluorine, neon, and sodium: 2
and here are O 2, F, Ne, and Na + : R (pm) Z N 3 146 7 O 2 140 8 F 133 9 Ne? 10 Na + 102 11 Mg 2+ 72 12 Al 3+ 53.5 13 3
Ionization Energy Neutral atoms are stable: they do not spontaneously lose electrons. The process X X + + e will always take energy, because the positively charged ion will always be attracted to the negatively charged electron. The (positive) energy input required is the ionization energy, usually measured in ev or in kj/mol, with 1 ev/atom = 96.49 kj/mol Elements that have multiple electrons will have multiple ionization energies: X(g) X + + e X + X 2+ + e X 16+ X 17+ + e I 1 = 1st ionization energy I 2 = 2nd ionization energy I 17 = 17th ionization energy I 1 I 2 I 3 I 4 I 5 I 6 I 7 Na [Ne]3s 1 3s 2p Mg [Ne]3s 2 3s 3s 2p Al [Ne]3s 2 3p 1 3p 3s 3s 2p Si [Ne]3s 2 3p 2 3p 3p 3s 3s 2p P [Ne]3s 2 3p 3 3p 3p 3p 3s 3s 2p S [Ne]3s 2 3p 4 3p 3p 3p 3p 3s 3s 2p Cl [Ne]3s 2 3p 5 3p 3p 3p 3p 3p 3s 3s Ar [Ne]3s 2 3p 6 3p 3p 3p 3p 3p 3p 3s 4
with energies in kj/mol: I 1 I 2 I 3 I 4 I 5 I 6 I 7 Na 495.8 4562.4 Mg 737.7 1450.7 7732.7 Al 577.5 1816.7 2744.8 11577.5 Si 786.5 1577.1 3231.6 4355.5 16090.6 P 1011.8 1907.5 2914.1 4963.6 6274.0 21267.4 S 999.6 2251.8 3357 4556.2 7004.3 8495.8 27107.4 Cl 1251.2 2297.7 3822 5158.6 6540 9362 11018.2 Ar 1520.6 2665.9 3931 5771 7238 8781.0 11995.3 Chemistry ionic or covalent involves valence electrons only. Periodic trends in first ionization energies: 1. Ionization energy increases as a shell fills. 2. There is generally a small dip in ionization energy when an electron is added after a filled subshell. 3. Adding a fourth electron to the p subshell produces a small dip in ionization energy. 5
4. Ionization energies tend to decrease moving down a column. 5. Putting it all together, ionization energies are lowest in the lower left of the periodic table (Rb and Cs are most easily ionized to Rb + and Cs + ) and highest in the upper right (He to He + ). Transition Metals and Lanthanides Electron Affinities The electron affinity is to anions what the ionization energy is to cations: X(g) + e X energy change = EA 1 6
In general, electron affinities: 1. become more negative from left to right on the periodic table, 2. are positive for noble gases, 3. are positive for elements with s 2 and d 10 configurations, 4. become more negative moving from heavier elements within a column to lighter though the first-row elements (BCNOF) are an exception to this. As with ionization energy, you can also have a second, third, fourth, etc. electron affinity; e.g. X + e X 2 energy change = EA 2 7
energy required to make cation Na Na + + e always positive energy required/gained to make anion Cl + e Cl sometimes negative, sometimes positive energy gained by bringing ions together Na + + Cl NaCl(s) always negative Electronegativity There are multiple electronegativity definitions to use. The simplest to understand (based on what we ve done so far) is the Mulliken definition: IP + EA ξ = 2 The electronegativity is, in many ways, a half-way point between ionization energy and electron affinity, and gives a general measurement of whether an element is going to pick up or give away an electron in any given compound. Electronegativity trends arise from those of ionization energy and electron affinity: 1. Electronegativity tends to increase from left to right (Li to F) and from bottom to to (Cs to Li, or I to F). 2. The non-metals in the upper right corner of the periodic table are elements with high electronegativity (ξ 2.2), and tend to gain electrons in chemical reactions they are oxidants. 3. Metals (and in particular, metals in the lower left corner) are elements with low electronegativity (ξ 1.8) and tend to lose electrons in chemical reactions they are reductants. 8