Lecture 5 Trends in Atomic Properties Energies and

Lecture 5 Trends in Atomic Properties

Energies and sizes -The first ionization energy (IE) of an atom (M) is the energy required to form the positive ion M+ M M+ + e. The IE value reflects the energy of the orbital from which the electron is removed, and so depends on the principal quantum number (n) and effective nuclear charge (Zeff)

IE= Z 2 eff. R/n 2 The average radius of an orbital depends on the same factors <r> = n 2 a 0/Zeff -Smaller orbitals generally have more tightly bound electrons with higher ionization energies.

Horizontal trends -Increasing nuclear charge is accompanied by correspondingly more electrons in neutral atoms. -Moving from left to right in the periodic table, the increase of nuclear charge has an effect that generally outweighs the screening from additional electrons. -Increasing Zeff leads to an increase of IE across each period, which is the most important single trend in the periodic table. At the same time, the atoms become smaller.

the first ionization energy of boron is smaller than beryllium, and the first ionization energy of oxygen is smaller than nitrogen. Ionization of boron removes an electron from a 2 p orbital, which is less tightly bound than the 2 s involved in lithium and beryllium. Thus the IE of B is slightly less than that of Be Be: [He] 2 s 2 B: [He] 2 s 2 2 p 1 For N and O: The electrons removed when nitrogen and oxygen are ionized both come from 2 p orbitals. But we have to think about Hund’s Rule: in oxygen, the orbital is slightly destabilized, as there is energy required to pair the electrons. Thus the electron in oxygen is lightly easier to remove than the one in nitrogen N: [He] 2 s 2 2 p 3 O: [He] 2 s 2 2 p 4

Ionization energies (IE) and electron affinities (EA) for the elements Li-Na.

Vertical trends -The IE generally decreases down each group of elements. - for hydrogen and the elements of group 1, all of which have the (ns)1 outer electron configuration. The main influence here is the increasing value of principal quantum number n. ). There is a substantial increase of nuclear charge between each element, and although extra inner shells are occupied, they do not provide perfect shielding. Thus, contrary to what is sometimes stated, effective nuclear charge increases down the group.

Ionization energies for elements with (ns)1 outer electron configurations.

In the resulting balance between increasing n and increasing Zeff the former generally dominates, as in group 1. There is, however, nothing inevitable about this, and there are occasions in later groups where Zeff increases sufficiently to cause an increase of IE between an element and the one below it.

- the group 11 elements Cu, Ag and Au, where an ns electron is also being ionized. The increase of IE along period 4 between K (Z=19) and Cu (Z=29) is caused by the extra nuclear charge of 10 protons, partly shielded by the 10 added 3 d electrons. A similar increase occurs between Rb and Ag in period 5. In period 6, however, the 4 f shell intervenes giving 14 additional elements and leading to a total increase of Z of 24 between Cs and Au. There is a much more substantial increase of IE therefore, and Au has a higher IE than Ag Similarly irregular trends in IE may have some influence on the chemistry of p-block elements Orbital radii also depend on n 2 and generally increase down each group.

- There is another interesting feature of vertical trends, arising also from the way in which the periodic table is filled. For orbitals of a given l there is a more significant change, both in IE and size, between the first and second periods involved than in subsequent cases. - for s orbitals, where the IE decreases much more from hydrogen (1 s) to lithium (2 s) than between the lower elements. Such a distinction is reflected in the chemical properties of group 1 elements, hydrogen being nonmetallic and the other elements metals). Similar - although less dramatic, differences are found with 2 p and 3 d. Thus period 2 p-block elements are in many ways different from those lower in the p block, and 3 d series elements distinct from those of the 4 d and 5 d

States of ionization - The successive energies required to create more highly charged ions, M 2+, M 3+ …are the second, third, …IEs. The values always increase with the degree of ionization. When electrons are removed from the same shell, the main effect is that with each successive ionization there is one less electron left to repel the others. The magnitude of the change therefore depends on the size of the orbital, as electrons in smaller orbitals are on average closer together and have more repulsion. Thus with Be (2 s)2 the first two IEs are 9. 3 and 18. 2 e. V, whereas with Ca (4 s)2 the values are 6. 1 and 11. 9 e. V, The third IE of both elements is very much higher (154 and 51 e. V, respectively) because now the outer shell is exhausted and more tightly bound inner shells (1 s and 3 p, respectively) are being ionized. The trends are important in understanding the stable valence states of elements.

Electron Affinity - Electron Affinity is the energy “required” to put an electron into an atom: M + e - M-This is an exothermic process. Energy is released upon the addition of an electron (in most cases); this will be a negative number.

• For most species except the noble gases (Group 18), alkaline earth metals (Group 2), and Group 12 (Zn, Cd, Hg) the process is exothermic (positive EA and energy is released) and stable ions are formed In general, and with a couple of exceptions, EA increases across a period

• Groups 2, 12, and 18 have negative EA values because the added electron must fit into a higher energy orbital and this requires expenditure of energy. Alkaline earths are all ns 2; the added electron must go – into the higher energy np orbital. Inert gases are all ns 2 np 6; the added electron must go – into the higher energy (n+1)s orbital. Group 12 are all ns 2 (n-1)d 10; the added electron • must go into the higher energy np orbital

* We know from the previous discussion on shielding that as we are filling a period, the added electrons do not shield a full nuclear charge and the dominant force is the increasing nuclear charge. As the nucleus becomes more highly charged it is able to stabilise added electrons more effectively. This is why the EA increases across a period.

• For each period, EA becomes smaller when moving from Group 14 to Group 15. By way of example, EA for C is larger than for N. Again this can be rationalised by considering where the electron goes: C: 1 s 2 2 p 2 N: 1 s 2 2 p 3

* Adding an electron to C allows it to go unpaired into the one empty 2 p orbital. Adding an electron to N however means that it must pair for the first time in a p orbital. This makes the process a little less favourable for the creation of N- and so EA is less than for the creation of C-.

Ionization energies (IE) and electron affinities (EA) for the elements Li-Na.