Chemistry 1000 Lecture 9: Periodic trends. Marc R. Roussel

Transcription

1 Chemistry 1000 Lecture 9: Periodic trends Marc R. Roussel

2 A qualitatively correct result of the Bohr theory Some results of the Bohr theory are qualitatively correct, even for multi-electron atoms, but need to be reinterpreted. Recall, from Bohr theory, rn n 2. We call a set of orbitals with the same n a shell. The average distance of the electrons from the nucleus is roughly proportional to n 2, so the Bohr theory is qualitatively correct. This average distance however increases slightly with l, which the Bohr theory doesn t predict.

3 Screening Suppose that we consider two shells with different n. The shell with the larger n has most of its electron density outside the shell with smaller n. Electrons in the small-n shell therefore screen the electrons in the large-n shell from the nucleus relatively effectively. Electrons in higher-n shells have an effect on lower-n electrons (because the high-n orbital does have some density inside the low-n orbital), but it s small. If there are other electrons in the same shell, these also have some screening effect on each other.

4 Effective nuclear charge The effective nuclear charge experienced by an electron takes into account the screening by other electrons. Effective nuclear charges can be calculated numerically using quantum chemical methods. The effective nuclear charge is slightly different for orbitals of different l values within a shell because of slightly different average distances to the nucleus. As we go across a period, the effective nuclear charge increases because the charge of the nucleus is increasing, but there is only partial screening by additional electrons added to the valence shell. The core electron configuration doesn t change across a period.

5 Effective nuclear charge Trend along a period 8 7 3s 3p 6 Z eff Na Mg Al Si P S Cl Ar

6 Atomic size Orbitals don t have a sharp cut-off. They are spongy. So how can we define an atomic size? Covalent radius: From compounds with single covalent bonds, whether network solids (e.g. diamond) or molecular compounds (e.g. F 2 ) For elements, the radius is half the distance between the nuclei. For compounds, the distance between the nuclei is the sum of the radii.

7 Atomic size Metallic radius: Same idea as the covalent radius, but in a metal. What s the difference, then? In a covalent compounds, some of the valence electrons are shared between two atoms. In a metallic compound, the valence electrons are shared throughout the metal. van der Waals radius: Mostly for noble gases, based on shortest distance between atoms in a solid crystal in which the atoms are not bonded to each other. Important note: These radii are not strictly comparable to each other, so we should always try to compare measurements of the same kind.

8 Atomic size Atomic size increases as we move down a group because we are adding shells and r n 2 /Z eff. In a period, for the main-group elements, Zeff increases as we move from left to right, so atomic size decreases.

9 Atomic size Trend down a group r/pm Li Na K Rb Cs

10 Atomic size Trend across a period r/pm Na Mg Al Si P S Cl

11 Ionic radius Measured analogously to covalent radius, but using crystals of an ionic compound. Bootstrapping problem: You need one radius in order to be able to assign the rest. Convention: r(o 2 ) = 1.40 Å Generally, similar trends observed as for atomic radii. Removing electrons (esp. if a shell is emptied) results in cations being smaller than the neutral atoms from which they are formed. The smallest ion in an isoelectronic series has the highest Z.

12 Ionic radius Example: Put the following ions in order of increasing size: O 2, F, Na +, Mg 2+ Answer: Mg 2+ (72 pm), Na + (102 pm), F (133 pm), O 2 (140 pm)

13 Ionization energy Ionization energy = 1 st ionization energy = I 1 Ionization energy decreases as we go down a group because the outer electrons are farther from the nucleus and thus more loosely held. As an overall trend, ionization energy increases as we move from left to right across a period because the effective nuclear charge increases and the size of the atom decreases, both of which increase the electrostatic force between the valence electrons and the nucleus.

14 Ionization energy (continued) We must also consider electronic configuration. Going from an ns 2 to an ns 2 np 1 configuration, the ionization energy decreases in the first few periods because the p orbital is higher in energy than the s. Going from an ns 2 np 3 to an ns 2 np 4 configuration, the ionization energy decreases because pairing electrons in an orbital increases electron-electron repulsion, thus making it easier to remove one.

15 Ionization energy Trend down a group I 1 /kj mol H Li Na K Rb Cs

16 Ionization energy Trend across a period for the 1st IE I 1 /kj mol Na Mg Al Si P S Cl Ar

17 Second ionization energy The second ionization energy is the energy required to remove a second electron from an atom, i.e. the energy required to remove an electron from X + forming X 2+. Second ionization energies can be understood using the same principles as first ionization energies, but you have to take into account the electron configuration of the singly ionized ion that is the starting point.

18 Second ionization energy Trend across a period for the 2nd IE I 2 /kj mol Na Mg Al Si P S Cl Ar

19 Enthalpy of electronic attraction This is called electron affinity in the textbook. ea H is the enthalpy change for the process A + e A in the gas phase. ea H is generally negative. A more negative value means a stronger attraction for electrons. Some elements have essentially no ability to accept an additional electron. Examples: Be, Mg, N, noble gases. Why?

20 Trend across a period ea H tends to become more negative as we go from left to right in a period because of increasing effective nuclear charge and decreasing atomic radius. Exceptions to the previous observation can generally be rationalized in terms of the electron configurations of the atom and anion.

21 Enthalpy of electronic attraction Trend across a period ea H/kJ mol Na Mg Al Si P S Cl Ar

22 Enthalpy of electronic attraction Group trends ea H tends to become less negative as we move down a group because of increasing atomic radius, but there are many exceptions. Second-period elements tend to have less negative ea values than the corresponding third-period element because of the small size of the former which concentrates the electrons in a small region of space and makes it energetically less favorable to add another electron. The trend is reversed in groups with low ea H values.

23 Enthalpy of electronic attraction Group 1 is normal ea H/kJ mol H Li Na K Rb

24 Enthalpy of electronic attraction Group 14 shows the off-trend second-period element ea H/kJ mol C Si Ge Sn Pb

25 Is there any special stability to noble-gas configurations? We often see claims that the propensity of atoms to form ions with noble-gas configurations is due to some special stability of the noble-gas electron configuration. If that were the case, it should be difficult to remove an electron from the F ion. The amount of heat (enthalpy) required to do so is ea H, which turns out to be 328 kj mol 1 for F. As ionization energies go, this is not a particularly large value. The context for studying issues of stability always has to be a balanced reaction in which we consider reactants and products. Avoid thinking in terms of any special stability properties of octets (even though that is often the end result in reactions involving main-group elements).

26 Electronegativity Usual symbol: χ An empirical measure of the tendency of an atom to attract electrons There are in fact several electronegativity scales. The Mulliken electronegativity scale is linearly related to I ea H. (Recall, ea H is negative if an electron can be added to an atom.) The more commonly used Pauling scale is based on bond dissociation energies, but correlates very well with the Mulliken scale. Electronegativity decreases as we go down a group. Electronegativity increases as we go from left to right across a period.

27 Electronegativity Trend down a group for χ χ H Li Na K Rb Cs Fr

28 Electronegativity Trend across a period for χ χ Na Mg Al Si P S Cl

29 Knowing the periodic table Important: You are expected to learn the names, symbols, and positions in the periodic table of all the elements in the first four periods (up to and including Kr).