Chapter 8: Periodic Properties

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# Chapter 8: Periodic Properties

## Chapter 8: Periodic Properties

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##### Presentation Transcript

1. Chapter 8: Periodic Properties • Nerve cells must pump Na+ ions out of the cell and K+ ions into the cell to complete the transmission of a signal. • Both are group 1A elements – yet the cells can distinguish them based on their size! • Size of atoms or ions is an example of a periodic property.

2. History • Elements like gold (Au) can be found in their pure state in nature and were known to the ancient humans. • Others are more rare and occur only in compounds, which until the 19th century were difficult to isolate. • In 1800, only 31 elements had been identified. By 1865, the number was 63. • www.ptable.com

3. History • In 1869, Dmitri Mendeleev proposed a classification scheme based on the fact that physical and chemical properties begin to repeat when the elements are arranged in increasing weight. • As many elements were still not yet discovered, Mendeleev left blank spaces where he predicted that these elements would be eventually isolated.

4. History

5. History • In 1913, Henry Moseley found that when an element is bombarded with high energy electrons it will produce an X-ray with a unique frequency. • He then arranged them in order of frequency and assigned them a whole number, which we now know as their atomic number.

6. Orbitals • Subshells are described by the combination of the n and l quantum numbers. • Energy increases for each subshell according to Aufbau filling order. • 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.

7. Electron Configurations • Ground state = lowest possible energy state for all of the electrons. • Each subshell in the Aufbau order fills completely before starting the next subshell. • Remember, s = 1 orbital, p = 3 orbitals, d = 5 orbitals, and d = 7 orbitals and EACH orbital can hold 2 electrons per quantum number rules.

8. Electron Configurations • Use a superscript after each subshell to indicate the number of electrons in that subshell. • H: 1s1 • He: 1s2 • Li: 1s2 2s1 • Be 1s2 2s2 • B: 1s2 2s2 2p1

9. Electron Configurations • Imagine doing a configuration for Sr or W or Pb!!! • Shorthand method • LEP #8. • Exceptions.

10. Subshells and the Periodic Table

11. Core and Valence Electrons • Core electrons are very close to the nucleus and are the completed shell(s) of electrons. • Valence electrons are the outermost shell of electrons. These are loosely held and provide for all of the bonding interactions. • For the main group elements, the number of valence electrons = the group number.

12. Orbital Diagrams • Hund’s Rule – when electrons are placed into degenerate orbitals, they will fill them singularly with parallel spins. • For outermost shells, use boxes, circles, or lines to represent each orbital. • Use a  for spin = +1/2 and use a  for spin = -1/2. • LEP #1

13. Periodic Properties • Elements in the same group tend to behave similarly. • Ex) Group 1A metals – Li, Na, K, Rb, Cs • Some, though, have major differences. • Ex) O and S • Electronic properties of valence electrons are always the same. • The chemical properties of the elements are largely determined by the number of valence electrons they contain.

14. Effective Nuclear Charge • The attraction of the electron to the nucleus depends on two factors. • Factor one is the magnitude of the charge. • Factor two is the distance between the charge. • In a many electron atom, each electron is attracted to the nucleus and repelled by the other electrons.

15. Effective Nuclear Charge • The core electrons versus the valence electrons. • Argon, 10 core and 8 valence electrons.

16. Effective Nuclear Charge • Core electrons “shield” the valence electrons from the nucleus. • Zeff= Number of Protons – Number of Core Electrons • Zefffor Na = +1, for Cl = +7 • Trend is: ________________________

17. Size of Atoms • Atoms are sometimes treated as hard spheres in models. • Defining atomic size is tricky, though, because atoms do not have defined boundary layers. • Non-bonding radii. • Bonding radii.

18. Size of Atoms • What trends do you see? • Any exceptions?

19. Electron Configurations of Ions • When a main group element gains or loses electrons, it does so with predictability. • Na (and other group 1A metals) always lose one electron. • Na: [Ne] 3s1 Na+: [Ne] + 1e- • Cl (and other group 7A halogens) always gain one electron. • Cl: [Ne] 3s2 3p5 + 1e-  Cl-: [Ne] 3s2 3p6

20. Electron Configurations of Ions • For transition metal ions, their configurations are NOT what is expected. • When filling, the 4s fills BEFORE the 3d orbital. • When removing electrons, though, they do not remove in reverse order! • ALWAYS remove the s electrons first. • Ex) Fe+2, Co+3 • Diamagnetic = no unpaired electrons, paramagnetic = unpaired electrons.

21. Size of Ions • What trends do you see?

22. Isoelectronic Series • Ions that contain the same number of electrons are said to be isoelectronic. • Place the following in order of increasing size: Na+, F-, Mg+2, O-2, and Al+3.

23. Ionization Energy (IE) • IE is the minimum amount of energy required to remove an electron from a gaseous atom or ion. • IE1 is the energy to remove the first electron from a neutral atom. • Ex) Na(g) Na+(g) + 1e- • IE2 is the energy to remove the next electron.

24. Ionization Energy • What trends do you see? • Any exceptions?

25. Successive IE’s • What do you see? • Why is it easier to remove the second electron from Cl than the second electron from Na?

26. Electron Affinity (EA) • EA is the amount of energy released (exothermic) when a gaseous atom gains an electron. • Ex) Cl(g) + e- Cl-1(g) ; DH = -349 kJ • An important note about sign and size of DH. • Some atoms will not accept an electron without having to absorb energy – these are said to have no measurable EA.

27. Electron Affinity • Ordered by groups. • Which group is the best? • Why? • Can we make predictions for transition metals?

28. Metals • Recall from Chapter 2 that metals: • Are usually solids at room temperature. • Are shiny and have luster • Are excellent conductors of heat and electricity. • Tend to lose electrons easily (form cations). • Metals oxides are said to be basic. • Metal oxide + water • Metal oxide + acid

29. Metals • Some metal ions can be identified by a flame test. • Seen below from left: Li, Na, and K.

30. Fireworks

31. Non-metals • Recall from Chapter 2 that non-metals: • Have a variety of appearances – some are gases, some are solids, and one is a liquid. Some are single atoms like metals, while others are diatomic or larger. • That are solids are brittle and poor conductors – except carbon! • Will either gain electrons (ionic) or share electrons (molecular). • Non-metal oxides are said to be acidic. • Non-metal oxide + water • Non-metal oxide + base

32. Metallic Character

33. Reactivity and Periodic Table • Groups 1A and 7A are highly reactive. • Group 1A loses 1 electron quite easily. • Thus, reactivity driven by I.E. • Reactivity increases from top to bottom within this group. • Why? • Group 7A gains 1 electron quite easily. • Thus, reactivity driven by E.A. • Reactivity greatest for F2 and Cl2 because they are gases.

34. Alkali metals and Halogens • The reaction of the alkali metals with water are all the same. • 2 Na(s) + 2 H2O(l)  2 NaOH(aq) + H2(g) + heat. • Reaction is highly exothermic. • The reaction of Al with the liquid Br2. • 2 Al(s) + 3 Br2(l)  2 AlBr3(s) + heat. • Reaction is also highly exothermic.

35. Alkali metals

36. Halogens

37. Groups 2A and 6A • Groups 2A and 6A are reactive, but less than the groups 1A and 7A. • Group 2A metals will still react with water, but much more slowly. • Group 6A non-metals (O2 and S) are similar.

38. Calcium metal

39. Why the differences? • Na  Na+ + 1e- ; H = +496 kJ • Ca  Ca+2 + 2e- ; H = +1,735 kJ • Similarly, if you compared energy to gain one electron for group 7A versus gaining two electrons for group 6A, then the latter is much harder.

40. Groups 3A, 4A, and 5A • Fairly unreactive with a few exceptions like Phosphorus. • Phosphorus comes in several forms including both white and red phosphorus. • White is the most unstable.

41. Red Phosphorus and Bromine

42. Group 8A • Group 8A forms no compounds at all with the exception of Xenon with Fluorine and Oxygen. • XeF2, XeF4, XeF6, XeO3, and XeO4 have been produced. • This is due to the relatively low I.E. of Xe.

43. Summary • Reactivity on the left-hand side of the periodic chart (metals!) is driven by the ease of losing electrons. • Reactivity on the right-hand side of the periodic chart (non-metals) is driven by the ease of gaining electrons.