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Special Topics for SOL 2 3 rd Power Point. Periodic Trends (Chap 14). Shorthand Electron Configurations. Shorthand configurations are a useful tool. Let’s look at an example for Y, Z=39 The electron configuration for yttrium is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 1

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shorthand electron configurations
Shorthand Electron Configurations
  • Shorthand configurations are a useful tool.
  • Let’s look at an example for Y, Z=39
  • 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(the bold and italics part)
  • That’s Kr and then we also just write whatever is left over.
  • [Kr] 5s24d1
you try
You Try…
  • Do a shorthand configuration for
  • Fe
  • Br
  • Rb
the answers
The Answers…
  • Do a shorthand configuration for
  • Fe = [Ar]4s23d6
  • Br = [Ar]4s23d104p5
  • Rb = [Kr]6s1
objective b http www rsc org chemsoc visualelements pages data intro groupvii data html
Objective Bhttp://www.rsc.org/chemsoc/visualelements/PAGES/data/intro_groupvii_data.html
  • Notice that the halogens all have an ending configuration of ns2np5. That means they have 7 valence electrons.
  • Similarly, alkali metal have 1 valence electron. Noble gases have 8, etc. All transition metals have 2.

F [He]2s22p5Cl [Ne]3s23p5

Br [Ar]3d104s2 4p5 I [Kr]4d105s2 5p5

At [Xe]4f14 5d106s2 6p5

objective b
Objective B
  • All of the transition metals have 2 valence electrons, with 2 exceptions. “d” electrons are not valence electrons. Why not?
  • Transition metals are where the d orbitals are being filled up. Here are the electron configurations for all of them.

Sc [Ar]3d14s2 Ti [Ar]3d24s2

V [Ar]3d34s2 Cr [Ar]3d54s1

Mn [Ar]3d54s2 Fe [Ar]3d64s2

Co [Ar]3d74s2 Ni [Ar]3d84s2

Cu [Ar]3d104s1Zn [Ar]3d104s2

objective b1
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.
  • That’s more stable than if Cr would have followed the pattern, and ended with “4s23d4”

Cr [Ar]3d54s1

objective b2
Objective B
  • 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.

Cu [Ar]3d104s1

objective b3
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.
objective c
Objective C
  • The periodic table allows you to predict trends in certain properties.
  • Get out a periodic table and put these trends as notes on your periodic table.
  • The first trend is Atomic radius.
  • Atomic radius is the size of the atom. It’s defined as ½ the distance between two nuclei which are bonded together.
objective c1
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.
  • A cation is always smaller than the atom it is formed from.
  • An anion is always larger than the atom it is formed from.
objective c http www chem1 com acad webtext atoms atpt images ionic radii jpg
Objective Chttp://www.chem1.com/acad/webtext/atoms/atpt-images/ionic_radii.jpg
  • Since cations lose electrons to form positive ions and anions gain electrons to form negative ions, it should make sense that they are SMALLER than the atom.
objective c2
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.
  • The energy required to remove the second electron is the second ionization energy. And so on…
  • Metals always have LOWER ionization energies than nonmetals.
  • That is because metals tend to lose electrons and nonmetals tend to gain them.
objective c3
Objective C
  • It is VERY MUCH easier to remove a valence electron (an electron in the highest energy level) than an “inner core” electron.
  • The inner core electrons are ANY electrons which are not VALENCE electrons.

Na = 1s22s22p63s1

White = inner core electrons and Blue = Valence electrons

objective c http www knowledgerush com wiki image 8 87 linuspauling jpeg
Objective Chttp://www.knowledgerush.com/wiki_image/8/87/LinusPauling.jpeg
  • Electronegativity is measured on a scale from 0.0 to 4.0.
  • By definition, F is the most electronegative element at 4.0.
  • Nonmetals have a high electronegativity.
  • Metals have a low electronegativity.
electronegativity
Electronegativity
  • Think of this as the “greediness” of an atom not only holding on to it’s own electrons, but ALSO wanting to “steal” electrons from other atoms.
the trends
The Trends
  • Atomic Radius AND Ionic Radius increase as you go down a group.
  • Atomic Radius AND Ionic Radius decrease as you go from left to right across a period.
  • Electronegativity AND Ionization Energy decrease as you go down a group.
  • Electronegativity AND Ionization Energy increase as you go from left to right across a period.

Note the trends are opposites. Draw some arrows on your periodic table to help you remember the trends.

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