Lecture 14: Periodicity of the elements (Ionization energy and electron affinity)

Ionization Energy

• Minimum energy needed to remove an electron from a gaseous atom.

• Endothermic process.

• Valence electron is the easiest to remove.

• $M(g) + IE_1 \rightarrow M^{1+}(g) + 1e^-$

• $M^{+1}(g) + IE_2 \rightarrow M^{2+}(g) + 1e^-$

• First ionization energy ($IE_1$) = energy to remove an electron from a neutral atom.

• Second ionization energy ($IE_2$) = energy to remove an electron from a +1 ion; and so on.

General Trends in First Ionization Energy

• The larger the effective nuclear charge on the electron, the more energy it takes to remove it.

• The farther the most probable distance the electron is from the nucleus, the less energy it takes to remove it.

• First ionization energy ($IE_1$) decreases down the group.

  • Valence electron is farther from the nucleus.

• First ionization energy ($IE_1$) generally increases across the period.

  • Effective nuclear charge increases.

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Trends in First Ionization Energy

• Decreasing ionization energy down a group and increasing ionization energy across a period. Specific values (kJ/mol) are given for elements in periods 1-6.

• Group 1A (Alkali metals) show the lowest ionization energies, while Group 8A (Noble gases) show the highest.

• Values are given for H (1312), Li (520), Na (496), K (419), Rb (403), Cs (376) in Group 1A.

• Values are given for He (2372), Ne (2081), Ar (1521), Kr (1351), Xe (1170), Rn (1037) in Group 8A.

• Other values are given for elements in Groups 2A through 7A, illustrating the general increasing trend across the period.

• There are some irregularities in the trend, which are explained later.

Example 2 – Choosing the Atom with Higher First Ionization Energy

• Al or S: S has a higher first ionization energy because it is further to the right on the periodic table.

• As or Sb: As has a higher first ionization energy because it is further up on the periodic table.

• N or Si: N has a higher first ionization energy because it is further up and to the right on the periodic table.

• O or Cl? This example involves opposing trends, and requires more nuanced consideration.

Irregularities in the Trend

• Ionization Energy generally increases from left to right across a Period EXCEPT from Group 2A to 3A, and from Group 5A to 6A.

• Examples used to illustrate the irregularities in the trend:

  • Be: $1s^2 2s^2$

  • B: $1s^2 2s^2 2p^1$

  • N: $1s^2 2s^2 2p^3$

  • O: $1s^2 2s^2 2p^4$

• It is easier to remove an electron from B than from Be because B loses an electron from the 2p subshell, while Be loses an electron from the full 2s subshell.

• It is easier to remove an electron from O than from N because O gains a half-full 2p subshell configuration upon ionization.

Irregularities in First Ionization Energy Trends

• Be: $1s^2 2s^2$

• B: $1s^2 2s^2 2p^1$

• $Be^+: 1s^2 2s^1$

  • To ionize Be, you must break up a full sublevel, which costs extra energy.

• $B^+: 1s^2 2s^2$

  • When you ionize B, you get a full sublevel, which costs less energy.

Irregularities in First Ionization Energy Trends

• N: $1s^2 2s^2 2p^3$

• O: $1s^2 2s^2 2p^4$

• $N^+: 1s^2 2s^2 2p^2$

  • To ionize N, you must break up a half-full sublevel, which costs extra energy.

• $O^+: 1s^2 2s^2 2p^3$

  • When you ionize O, you get a half-full sublevel, which costs less energy.

Trends in Successive Ionization Energies

• Removal of each successive electron costs more energy.

  • Shrinkage in size due to having more protons than electrons.

  • Outer electrons closer to the nucleus, therefore harder to remove.

• Regular increase in energy for each successive valence electron.

• Large increase in energy when starting to remove core electrons.

Successive Values of Ionization Energies for the Elements Sodium through Argon (kJ/mol)

• Table showing ionization energies (IE1 through IE7) for elements Na through Ar.

• For Na:

  • $IE_1 = 496$

  • $IE_2 = 4560$

• For Mg:

  • $IE_1 = 738$

  • $IE_2 = 1450$

  • $IE_3 = 7730$

• For Al:

  • $IE_1 = 578$

  • $IE_2 = 1820$

  • $IE_3 = 2750$

  • $IE_4 = 11600$

• For Si:

  • $IE_1 = 786$

  • $IE_2 = 1580$

  • $IE_3 = 3230$

  • $IE_4 = 4360$

  • $IE_5 = 16100$

• For P:

  • $IE_1 = 1012$

  • $IE_2 = 1900$

  • $IE_3 = 2910$

  • $IE_4 = 4960$

  • $IE_5 = 6270$

  • $IE_6 = 22200$

• For S:

  • $IE_1 = 1000$

  • $IE_2 = 2250$

  • $IE_3 = 3360$

  • $IE_4 = 4560$

  • $IE_5 = 7010$

  • $IE_6 = 8500$

  • $IE_7 = 27100$

• For Cl:

  • $IE_1 = 1251$

  • $IE_2 = 2300$

  • $IE_3 = 3820$

  • $IE_4 = 5160$

  • $IE_5 = 6540$

  • $IE_6 = 9460$

  • $IE_7 = 11000$

• For Ar:

  • $IE_1 = 1521$

  • $IE_2 = 2670$

  • $IE_3 = 3930$

  • $IE_4 = 5770$

  • $IE_5 = 7240$

  • $IE_6 = 8780$

  • $IE_7 = 12000$

• The data shows a regular increase in energy for each successive valence electron and a large increase in energy when starting to remove core electrons.

Trends in Electron Affinity

• Energy released when a gaseous neutral atom gains an electron.

• $M(g) + 1e^- \rightarrow M^{-1}(g) + EA$

• Defined as exothermic (-), but may actually be endothermic (+).

  • Alkaline earth metals & noble gases are endothermic.

• The more energy released (more negative), the larger the EA.

• Generally increases across a period.

  • Becomes more negative from left to right.

  • Lowest EA in group = alkaline earth metal or noble gas.

  • Highest EA in group = halogen.

Electron Affinities (kJ/mol)

• Values are provided for elements in periods 1-6. Note that >0 indicates an endothermic process.

• Group 1A values are negative, indicating exothermic processes, while Group 2A values are positive.

• Specific values include:

  • H: -73

  • Li: -60

  • Na: -53

  • K: -48

  • Rb: -47

• Halogens (Group 7A) have high negative electron affinities:

  • F: -328

  • Cl: -349

  • Br: -325

  • I: -295

• Noble Gases (Group 8A) have positive electron affinities (>0).