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Chapter 14

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Chapter 14

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  1. Chapter 14 Chemical Periodicity

  2. Objective A • Chapter 14 is a very short chapter. We also already know some of what is in this chapter. • The Periodic Table groups elements according to their properties. • Look at the first group. It has H and Li and Na, etc. • All of these elements behave in the same ways. • If you know the properties of Li and Na, you can make a very good inference that K and Cs will behave the same way.

  3. Objective A • You can only say that about elements in the same group (column, going up and down) • Elements in the same period can have different properties, so you can’t make the assumption that Na will behave like Mg or Mg will behave like Al. • We can use the electron configuration to make assumptions about the element. Elements in the same group will have the same “ending” to their electron configuration. Li [He]2s1Na [Ne]3s1 K [Ar]4s1 Rb [Kr]5s1 Cs [Xe]6s1 Fr [Rn]7s1

  4. Objective B • Alkali metals all have 1 electron in their highest occupied energy level. • Noble gases all have 8 electrons in their highest occupied energy level. That’s not true for Helium, but remember that Helium only has the 1s orbital (so when it has 2 electrons, it’s highest occupied energy level is full). • We call the electrons in the highest occupied energy level the “valence electrons.” He 1s2Ne [He]2s22p6 Ar [Ne]3s23p6 Kr [Ar]3d104s24p6 Xe [Kr]4d105s25p6 Rn [Xe]4f145d106s26p6 Shorthand configurations…we’ll learn about those in just a sec…

  5. Objective B • Notice that the halogens all have an ending configuration of ns2np5. That means they have 7 valence electrons. • Let’s also look at the transition metals. We’ll only look at the first row, called the first transition series of elements. F [He]2s22p5Cl [Ne]3s23p5 Br [Ar]3d104s2 4p5 I [Kr]4d105s2 5p5 At [Xe]4f14 5d106s2 6p5

  6. Objective B • All of the transition metals have 2 valence electrons, with 2 exceptions. • Transition metals are where the d orbitals are being filled up. Here are the electron configurations for all of them.

  7. Objective B • Notice that Cr and Cu are “exceptions.” • They both have 1 valence electron. They do this because in the case of Cr, moving an electron from the 4s level to the 3d level gives us a half full set of d orbitals.

  8. Objective B • That’s more stable than if Cr would have followed the pattern, and ended with “4s23d4” • Similarly, Cu has 1 electron in the 4s energy level and 10 in the 3d level, because having a full set of d electrons is also more stable.

  9. Objective B • We won’t do much in this class with the second transition series or the third. • We will talk about some of those elements, but most of the patterns for the first transition series will hold true for the others.

  10. Objective B • The “inner transition metals” are the lanthanide and actinide series. • That’s where the f electrons are filled up. • That’s about all I’m going to say about that, except that Glenn Seaborg was the first to propose the existence of the actinides. • He called it the Actinide Hypothesis, and many scientists felt that he was wrong. However, he was proven to be correct.

  11. Objective B • Shorthand configurations are a useful tool. • As you can see, when you get a lot of electrons, the configuration can get pretty long. • Let’s look at an example for Y, Z=39

  12. Objective B • The electron configuration for yttrium is • 1s22s22p63s23p64s23d104p65s24d1 • To do a shorthand configuration, we use the noble gas preceding the element and we put that in brackets.

  13. Objective B • 1s22s22p63s23p64s23d104p65s24d1 • The noble gas that precedes Y is Kr. • Kr electron configuration is 1s22s22p63s23p64s23d104p6, which we represent as [Kr].

  14. Objective B • 1s22s22p63s23p64s23d104p65s24d1 • I can replace the underlined part with [Kr], leaving me with a shorthand configuration of • [Kr]5s24d1

  15. Objective B • Do a shorthand configuration for • Fe • Br • Hg don’t forget that after 6s comes 4f and 5d!

  16. Objective B • Do a shorthand configuration for • Fe = [Ar]4s23d6 • Br = [Ar]4s23d104p5 • Hg = [Xe]6s24f145d10

  17. Objective C • The periodic table allows you to predict trends in certain properties. • Atomic radius is one of those properties. • Atomic radius is the size of the atom. It’s defined as ½ the distance between two nuclei which are bonded together.

  18. Objective C • Ionic radius is another property • It is the size of an ion. Ionic radius is fairly similar to atomic radius. • A positive ion is also called a CATION. • A negative ion is also called an ANION.

  19. Objective C • A cation is always smaller than the atom it is formed from. • An anion is always larger than the atom it is formed from. • Since cations lose electrons to form positive ions and anions gain electrons to form negative ions, that should make sense.

  20. Objective C • Ionization energy is the amount of energy required to remove an electron from a gaseous atom. • The energy required to remove the first electron is called the FIRST IONIZATION ENERGY.

  21. Objective C • The energy required to remove the second electron is the second ionization energy. • Metals always have LOWER ionization energies than nonmetals. • That is because metals tend to lose electrons and nonmetals tend to gain them.

  22. Objective C • It is easier to remove a valence electron (an electron in the highest energy level) than an “inner core” electron. • The inner core electrons are the electrons in the lower energy levels. For example, sodium has 1 valence electron in the 3rd energy level. Sodium has 8 electrons in the 2nd energy level and 2 in the 1st. The 10 electrons in the 1st and 2nd levels are called “inner core” electrons.

  23. Objective C • Electronegativity is the tendency of an element to attract electrons to itself when they are bonded to another element. • Nonmetals have a very high electronegativity and metals have a very low electronegativity.

  24. Objective C • Electronegativity is measured on a scale from 0.0 to 4.0. • By definition, F is the most electronegative element at 4.0.

  25. Objective C • Now the trends….. • Atomic radius increases as you go down a group. • Atomic radius decreases as you go from left to right across a period.

  26. Objective C • Now the trends….. • Ionic radius increases as you go down a group. • Ionic radius decreases as you go from left to right across a period. • However, there is a big jump in size between groups 4A and 5A (this is where you switch from cations to anions).

  27. Objective C • Now the trends….. • Ionization energy decreases as you go down a group. • Ionization energy increases as you go from left to right across a period.

  28. Objective C • Now the trends….. • Electronegativity decreases as you go down a group. • Electronegativity increases as you go from left to right across a period.

  29. Objective C • If it helps…put arrow on one of your periodic tables showing the trends.

  30. The End