Students will be able to:
compare the relative sizes of the atoms of elements based upon their location on the periodic table and explain why & how the radii of atoms changes across a period and down a group
identify the property of ionization energy and discuss and periodic trend of ionization energy
identify the property of electron affinity and discuss and explain the periodic trend of electron affinity
describe the effects of atomic radius, ionization energy and electron affinity on the observed properties of atoms; e.g., formation of ions and reactivity
Periodic Table of the Elements
compare the relative sizes of the atoms of elements based upon their location on the periodic table and explain why & how the radii of atoms changes across a period and down a group
identify the property of ionization energy and discuss and explain the
identify the property of electron affinity and discuss and explain the
describe the effects of atomic radius, ionization energy and electron affinity ties of atoms; e.g., formation of ions and reactivity
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36 This booklet is simply a quick review of the fundamental trends of atomic radius, ionization energy, electron affinity and general atomic properties as they can be seen using the periodic table. Much more on the periodic table will be incorporated into additional units of study.
Part 1: ATOMIC RADIUS
Atomic radius is most closely aligned with the effective nuclear charge (Zeff) experienced by the electrons on an atom. The effective nuclear charge is a measure of the nuclear charge experienced by an electron. As the effective nuclear charge increases for electrons, they will be pulled closer to the nucleus. For example, the nuclear charge on nitrogen is 7+, while for oxygen it is 8+. Thus, the 2s electrons on the two atoms experience different attraction by the nucleus – the 2s electrons of oxygen are more strongly attracted, which leads to a smaller radius for oxygen compared to nitrogen.
Atomic radius generally decreases across the periodic table as the number of protons increases and the effective nuclear charge increases.
Atomic radius generally increases down a group of elements as the effective nuclear charge decreases on the electrons in larger shells and the repulsion between electrons increases.
Two atomic radius measurements are important in our discussion: bonding radius and nonbonding radius.
Figure 9. The trend of atomic radius can easily be seen in this figure - the increasing number of protons left-to-right and the increasing electron shells top-to-bottom cause this trend.
Figure 10. In graph form, the trend of atomic radius across a period is quite apparent.
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The radius of atoms changes when they bond because the bonded atoms’ nuclei are pulling on one another’s electrons – this pulling causes the atoms in the bond to become distorted and “squeezed” by the electrostatic attraction between the nuclei and electrons of the atoms involved.
The nonbonding radius of atoms is somewhat larger than the bonding radius because the atoms are not subjected to the pulling force of another atom’s nucleus.
Atomic radius is measured in the SI unit of picometer, pm, which is 10-12 meter, or 1/1 000 000 000 000 meter. Atoms have atomic radii ranging from about 30 pm for hydrogen to about 300 pm for larger metals in Groups 6 and 7.
Part 2: IONIZATION ENERGY
As atoms become smaller, you should be able to imagine that it would require more energy to remove electrons – the attraction is greater between the nucleus of a small atom and the atom’s electrons than between the nucleus of a larger atom and its electrons. Thus, ionization energy – the amount of energy required to remove an electron from a gaseous atom or ion – increases as atoms decrease in atomic radius and decreases as atoms become larger.
For example, the ionization of the first electron from sodium(g) is shown here:
Na(g) → Na+(g) + e- ∆H = +496 kJ/mol
This energy is called the first ionization energy, I1, of sodium because it is the energy required to remove the most loosely-held valence electron. The second ionization energy, I2,of sodium is the amount of energy required to remove the most loosely-held electron from the sodium ion, Na+ (i.e., the second electron). Notice the very large (almost ten-fold) increase in the ionization energy when the second electron is removed, an observation for which we will explore an explanation soon:
Na+ (g) → Na2+(g) + e- ∆H= +4560 kJ/mol Figure 11. The trend of ionization energy is explained by evaluating the size of the
atoms. You should be able to explain several small but notable exceptions.
38 Table 1. The removal of valence electrons requires a significantly lower energy than does the removal of non-valence electrons, as is shown in this table of the first through seventh ionization energies for the Period 3 elements.
IE Na Mg Al Si P S Cl Ar
First 496 738 578 787 1012 1000 1251 1520
Second 4562 1451 1817 1577 1903 2251 2297 2665
Third 6912 7733 2745 3231 2912 3361 3822 3931
Fourth 9543 10540 11575 4356 4956 4564 5158 5770
Fifth 13353 13630 14830 16091 6273 7013 6540 7238
Sixth 16610 17995 18376 19784 22233 8495 9458 8781
Seventh 20114 21703 23293 23783 25397 27106 11020 11995
The trends in ionization energy can be readily explained by looking at effective nuclear charge: electrons farther from the nucleus are more shielded (that is, experience a lesser nuclear charge) than closer-in electrons. This shielding is caused by the number of repulsive electrons between any given electron and the nucleus. Clearly, for example, s electrons are less shielded than are the p electrons in the same sublevel. Thus, the p electrons exhibit lower ionization energy than their same-level s electrons – and they certainly require less energy to remove than any (n – 1) electrons.
The trend of ionization energy is influenced by the size of the atom. It seems intuitive that a smaller atom exerts a greater attraction between the nucleus and electrons than does a larger atom. Thus, the smaller atoms exhibit higher ionization energies than the larger atoms.
Although the general trend of ionization energy is an increase left-to-right across the periodic table, there are cases where a decrease in ionization energy is evidenced. For example, the small decreases noticed as we move from Group 15 to Group 16 can be explained in terms of Hund’s rule: the atoms of Group 16 possess filled p-orbitals, which increases the electron repulsions and makes the first ionization energy lower than is seen in the ½-filled p-orbitals of Group 15 where the electron-electron repulsion is not as great because no electrons are paired.
Figure 12. For the same atom, notice that the outermost electron would experience a lesser effective nuclear charge than a closer-in electron. Thus, the ionization energy of the first electron is less than that of the second or subsequent electrons.
39 Part 3: ELECTRON AFFINITY
Ionization energy expresses the amount of energy required to remove electrons from gaseous atoms to form positively-charged ions. It is also expected to see atoms gain electrons to fill their valence shells. A measure of the energy change associated with the formation of negatively-charged ions when gaseous atoms take electrons into their valence shells is expressed as the electron affinity of atoms. The addition of electrons to atoms is generally exothermic, which results in a negative sign on the energy change1, as shown here:
Cl(g) + e- → Cl–(g) ∆H= –349 kJ/mol
The more negative the value of ∆H – corresponding to a larger release of energy – the greater the electron affinity of an atom; this generally means that the negative ion is more stable relative to the free atom. As ∆H becomes less negative there is lower tendency for the negative ion to form. Indeed, a positive ∆H indicates that the addition of energy is required for the addition of an electron.
Practice 2.1 Place each set in order of increasing atomic radius:
P, Se, S, As Na, Be, Mg
Arrange the following in increasing order of predicted first ionization energy: Ne, Na, P, Ar, K
1There are at least two widespread sign conventions on ∆H for electron affinity. Recognize that the bottom-line is that the addition of an electron is favorable for a nonmetal atom, while it is less favorable for a metal atom or a nonmetal ion that is already charged to a full valence.
Figure 13. The electron affinity of the atoms generally becomes more negative (increases) across a period.
40 Consider the ionization energies seen here:
I1 = 577.5 kJ, I2 = 1816.7 kJ, I3 = 2744.8 kJ, I4 = 11577 kJ, I5 = 14842 kJ, I6 = 18379 kJ
In what group is this atom likely located? Justify your response.
Place the following in predicted order of electron affinity from greatest electron affinity (most negative) to lowest electron affinity (most positive).
Na, Ar, Cl, Si
The Group 1 atoms have slightly negative values of electron affinity, while the Group 2 metals have slightly positive values. Explain this observation.
Why might the Group 15 elements have more positive electron affinity values than other nonmetals, even though we know they typically form negative ions?
ADVANCED PLACEMENT CHEMISTRY