From last lecture...

1. From last lecture... How do we explain the periodic and group properties of the elements? First, we need to know that chemical reactions involve the electrons of an atom only. The nucleus is not involved in chemical processes!

2. Electrons are weird: they act like matter = have mass, they act like energy = show wave properties. Very small, fast-moving subatomic particles do not behave like ordinary objects in motion.

3. Opposite charges attract each other. Nucleus of atom has + charge from protons.Electrons outside nucleus have - charge. So, electrons are pulled towards an atom’s nucleus, but : electrons don’t fall into or touch nucleus.

4. Energy is required to keep + and - charged particles separated. Therefore, the electrons in an atom have energy values. All of an atom’s electrons are not at same distance from nucleus. => Not all electrons have the same energy.

5. Electron energy levels in atoms have discrete values. This is like saying you can live at certain fixed distances from school: 1 mile, 1.5 miles, 2 miles, 2.9 miles... but NOT at any other distance (say 1.3 or 2.5 miles). => Electron energy values are quantized (= fixed, constant).

6. Quantum Mechanics is the branch of chemistry/physics that deals with the behavior of electrons. Complex equations have been developed to describe the energy of an electron in an atom. Electrons are moving and smaller than you can imagine.

7. The uncertainty principle tells us that we can not know both the energy and position of an electron at the same time. An orbital is an equation which describes a volume of space a certain distance from the nucleus in which there is about a 95% probability of finding an electron of a particular energy.

8. We use four quantum numbers to describe the state of an electron in an atom: n, principal quantum # tells how far the electron is from the nucleus; l, subshell quantum # describes the shape of the space in which the e- can be found.

9. m, tells how many orbitals of a kind there are; s, tells the spin of the electron.

10. n takes number values = 1,2,3...integer • think of n as a floor number: the higher the number value, the farther the electron is from the nucleus & the higher its energy.

11. l is assigned letter values = s, p, d, f • s-type orbitals are spherical, with nucleus at center; • p-type orbitals are shaped like 3-dimensional figure 8s, with nucleus at the cross-over; • d-type andf-type orbitals have multi-lobed shapes.

12. When we name an orbital describing the energy of an electron, we state both the n and l quantum numbers: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f All the orbitals of the same n value are described as a shell. All the orbitals of same n & same l value are termed a subshell.

13. The number of each kind of orbitals is given by the m quantum number. In any shell (= any value of n) there can be only: one s-type orbital; threep-type orbitals; five d-type orbitals; and seven f-type orbitals.

14. Remember: an orbital is a mathematical description of the energy of an electron, not a physical object like a box or jar. When we say an electron is “in” orbital ___, we mean the electron is in an energy state described by that orbital.

15. Each orbital can describe at most 2 electrons. • there can 0, 1 or 2 electrons “in” that orbital. Electrons spin on their axis like the earth, but either east west or west  east. When there are two electrons in the same orbital, one spins in each direction.

16. We say these electrons are “spin paired” and “in the same orbital.” Since there are only 2 possible values for s, two electrons can have the same n, l and m values only if they have opposite spins.

17. Electronic Configuration • lists orbitals according to their increasing energy, and •tells how many electrons are “in” or “occupy” each orbital. The relative energy levels and number of electrons in each orbital is found by experiment.

18. The highest energy level electrons in any atom are called the valence level electrons. The periodic table on next page is coded to show what type of orbitals are used by the valence electrons in each element. (f type orbitals are omitted for simplicity).

20. In describing the electronic configuration of an atom, we recognize that each element has one more proton and one more neutron than the one before it in the periodic table. We can pretend to “build up” an atom by packing more protons into the nucleus.

21. We show how the electrons are arranged outside the nucleus by describing the electronic configuration in terms of the: n (shell) quantum number, l (subshell) quantum number, and the number of electrons (1 or 2) “in” that orbital.

22. To write the electronic configuration of any element, we pretend to “build” the atom starting from hydrogen, atomic number 1. H is 1s1. 1 for n = 1 (period 1) S for spherical orbital. Superscript 1 tells there is 1 electron.

23. Li, atomic number 3= 1s2, 2s1 Nitrogen N, atomic # = 7 is configured as 1s2, 2s2, 2p3. Magnesium Mg, atomic # = 12, has the configuration 1s2, 2s2, 2p6, 3s2.

24. Check: what is the electronic configuration for Chlorine?

25. Chlorine Cl, atomic # = 17 is 1s2, 2s2, 2p6, 3s2, 3p5.

26. Shorthand notation We write the symbol of the last Nobel Gas to represent all the filled orbitals up to that point: Mg with 12 electrons long = 1s2, 2s2, 2p6, 3s2. Shorthand = [Ne] 3s2.

27. An element’s position on the Periodic Table reflects its electronic configuration. period number = quantum number n for highest energy electrons (the valence level). group number is related to the number of electrons in level n.

28. Elements in the same Group have the same valence level configuration, just at different levels n. This is why their chemical properties turn out to be similar.

29. Examples The elements of Group 1 all have a single valence level electron: H at 1s1; Li at [He] 2s1; Na at [Ne] 3s1; K at [Ar] 4s1; Rb at [Kr] 5s1; Cs at [Xe] 6s1; Fr at [Rn] 7s1.

30. Notice that all the Group 1 metals have 1 valence electron in an s orbital. What is different is how far out the electron is from the nucleus. (Larger n value, farther out: the valence e- for Li is closer to the nucleus than for K.) Also: all these metals have reactions involving just 1 electron.

31. More Examples The elements of Group 13 all have 3 valence level electrons: B is [He] 2s2, 2p1; Al is [Ne] 3s2, 3p1; Ga is [Ar] 4s2, 4p1; In is [Kr] 5s2, 5p1.