Periodic Properties

1. Periodic Properties

2. You learned how to write the electron configurations for ions. • Just remember that the valence electrons get added to or removed first! • So you always add to or remove from the highest principal energy level, or n value. • Write the electron configuration for Cl-, Mn2+ Ionic Electron Configurations

3. Can you see why metals tend to lose electrons? • Can you see why nonmetals tend to gain electrons? • What kind of e- configuration are they trying to achieve? • How many valence s and p electrons is this? • This is the basis for the Octet Rule: Elements in the Main Group tend to gain, lose, or share electrons in chemical rxns in order to have 8 valence electrons, or filled s and p subshells like the Noble Gases. Ionic Electron Configurations

4. You learned the Periodic Trend for atomic radii, that is the radii for neutral atoms. • What happens to the size of an atom when it loses electrons to become a cation? • Why does it become smaller? (think about Zeff and the shell that it loses e- from) • Likewise, what happens to an O atom when it gains 2 e- to become the oxide anion? • Why does it get larger? Ionic Radii

6. You know that an electron in an orbital may be promoted to a higher energy level orbital if it absorbs enough energy. • But you also know that electrons may be removed from an atom entirely. • Although it is typically the valence electrons which are removed, we can actually remove core electrons as well. • So if we put in enough energy, we can remove an electron (or more than one) entirely. Ionization Energy

7. The Ionization Energy is the energy required to remove an electron from an atom (or ion) in the gaseous state. • Li(g)→Li+(g) + e- • The first Ionization Energy, IE1 or Ei1, is the energy required to remove the outermost valence electron from a neutral atom. • The second IE, IE2, is the energy required to remove the second valence electron from the +1 cation. • What’s the IE3? Ionization Energy

8. How does IE vary down a group or across a Period? • Going down a group, you add a shell, so the valence electrons are further and further from the nucleus. As the atom gets larger, the valence electron is further away, and is loosely held. So it is easier to remove, and the IE is smaller. So IE decreases going down a group. • Going across a Period, the atomic size isdecreasing due to an increasing Zeff. Both of these means that the valence electrons are more tightly held, so the IE increases going across a period from left to right. Ionization Energy

9. Ionization Energy

10. Look at the following graphs, do you see any exceptions, or something you think is unusual? Ionization Energy

11. Ionization Energy

12. Ionization Energy

13. There are exceptions to the general trends, but many of them are explained easily. • For example, the Alkaline Earth metals have higher IE1 values than expected, higher than Group 13 values. • IE1 for Group 2 is removing an s electron, while IE1 for Group 13 is removing a p electron in the same energy level. • s electrons are slightly closer on average to the nucleus, are more tightly held to the atom, and so are harder to remove than a p electron on the same level. Ionization Energy

14. Also, Group 16 tend to have lower than expected IE values, and are often lower than the values for Group 15. • This is due to electron-electron repulsion effects. • Group 16 has a valence e- configuration of ns2np4, while Group 15 has a valence e- configuration of ns2np3. • Group 16 has the last valence electron doubled up in a p orbital, so there is higher electron-electron repulsions in this filled p orbital, so it is easier to remove an electron. • Group 15, on the other hand, has no filled p orbitals, so the electron-electron repulsions are lower, so it is harder to remove an electron from these p orbitals. • Group 15 has the p orbitals half-filled! More stable! Ionization Energy

15. Ionization Energy

16. Note that the IE values for the transition metals don’t increase very much (d-block) going across a Period. • Remember that the ns2 electrons are the valence electrons for the transition metals, so they are removed first, NOT the d electrons. Ionization Energy

17. What about higher IE values? How do you think they compare to the IE1 values? • For example, for Mg: how many valence electrons does Mg have? • Is there a trend for IE1, IE2, IE3 for Mg? • Do you think that there will be a large increase in IE for one of these? Ionization Energy

18. Mg has 2 valence electrons, 3s2. • When we remove the first 3s electron, we put in 738 kJ/mol and produce the Mg+ ion. • Do you think that it will be harder or easier to remove a second electron from a Mg+ electron? • When we remove the second 3s electron, we must put in 1451 kJ/mol (or roughly double) and we produce the Mg2+ cation. So the more positive the charge on an ion, the harder it is to remove an electron. This should make sense as cations are smaller than the neutral atom. Ionization Energy

19. What about the IE3 for Mg? Now what valence electron are we trying to remove? We have removed the 3s electrons, so we are now trying to remove a 2p electron! This is a core electron. • As we are now removing a core electron from a lower energy level (closer to nucleus) as well as removing an electron from a cation, the IE jumps sharply (roughly 8 to 10 times greater). It now requires 7733 kJ/mol to remove the third electron from Mg. Ionization Energy

20. We can relate this to the valence electrons of Group 2. They have 2 valence electrons, so they may be removed fairly easily, but once the valence electrons are removed, there is a huge increase in the next IE as we are now removing core electrons. • So you can always tell what group an atom is in given the successive IE values, as they are giving you the number of valence electrons! Ionization Energy

21. Ionization Energy

22. Example: An atom in Period 2 showed the following IE values: IE1 = 750 kJ/mol; IE2 = 1500 kJ/mol; IE3 = 2800 kJ/mol; IE4 = 5000 kJ/mol; IE5 = 38,000 kJ/mol; IE6 = 47,000 kJ/mol. What group is this atom a member of? What element is it? Ionization Energy

23. Ionization Energy

24. Electron Affinity, EA or Eea, is defined as the energy change when an electron is added to an atom (or ion) in the gaseous state: • F + e-→F- • How has the electron configuration of F changed once an electron is added? Can you see why nonmetals tend to gain electrons? • The electron affinity value gives us an idea of the attraction that atom has for an electron. Electron Affinity

25. For most atoms, energy is released when an electron is added so the EA is negative. • So the more negative the EA, the more likely the atom will gain an electron, and the resulting anion is more stable. • The Halogens have very negative EA values which should make perfect sense as they only need 1 electron to achieve the same electron configuration as a Noble Gas. • However, for the Noble Gases and a few other atoms (particularly Group 2 with their filled s subshell and to a lesser extent Group 15 with their half-filled p subshell), the EA is actually positive. This means that it is very difficult to add electrons to these atoms. Electron Affinity

26. Electron Affinities

27. You can add more than 1 electron to an atom, but the second electron is added to the anion: • O + e-→O- EA1 = -141 kJ/mol • O- + e-→O2- EA2 = 878 kJ/mol • As you might expect (and as is similar to the successive IE values), it is harder to add an electron to an anion as it has a lower Zeff, it is larger, and electron-electronrepulsions are high. Electron Affinity

28. However, oxygen and other nonmetals may gain more than 1 electron, but this is one of the reasons why it is very hard for nonmetals to gain 3 or more electrons. • For example, although N does form ionic compounds as nitrides where the N has gained 3 electrons (N3-), it forms many more covalent compounds. Electron Affinity

29. How does the EA vary going down a Group and across a Period from left to right? • First of all, the EA is the Periodic Property which has the most variation from the general trend. Electron Affinity

30. That said, IN GENERAL, the EA becomes less negative as you go down a group. • However, the change is small and is not always followed. • This is because the atomic radius is larger down the group which would tend to make the EA less negative, but the electron-electron repulsions are smaller in the larger atom, which tends to make EA more negative! • The EA generally becomes more negative across a Period from left to right (or going towards the Halogens). Electron Affinity

31. Electron Affinities

32. The Periodic Table is also roughly arranged by metallic character. • Remember the general properties of metals (review them)? • The more an element has these properties, the more metallic it is. • Metals tend to have lower IE values and less negative EA values as they tend to lose electrons to become cations. • Where is this on the Periodic Table? Metallicity

33. Based on the above, you should be able to predict that as you go down a Group, the metallicity increases. • Why? The atomic radius increases, so the IE decreases, and the EA becomes less negative (roughly). So it is easier to remove an electron, so it is more metallic! • What about going across a Period from left to right? • Now the atomic radius is decreasing, the Zeff is increasing, so the IE increases, and the EA becomes more negative (roughly). So the metallicity decreases across a Period from left to right. Metallicity

34. Metallicity

35. So what’s the most metallic element on the Periodic Table? (Not that anyone works with it, it is too radioactive and toxic!) Metallicity