The Periodic Table

1
The Periodic Table

• Jedediah Mephistophles Soltmann

2
Dmitri Mendeleev

• Studied the properties of elements and organized
  the elements by similar properties (families) and
  by increasing atomic mass.
• He left blanks for elements he knew had to exist,
  such as

3
Ekaaluminum (gallium)

• In 1871 Mendeleev predicted the existence of yet
  undiscovered element he named eka-aluminum
  (because of its proximity to aluminum in the
  periodic table). The table below compares the
  qualities of the element predicted by Mendeleev
  with actual characteristics of Gallium
  (discovered in 1875).

Property Ekaaluminum Gallium
atomic mass 68 69.3
density (g/cm³) 5.9 5.93
melting point (C) Low 30.15
oxide's formula Ea2O3 Ga2O3
chloride's formula Ea2Cl6 Ga2Cl6
4
Ekasilicon (Germanium)
Property Ekasilicon Germanium
atomic mass 72 72.59
density (g/cm³) 5.5 5.35
melting point (C) high 947
color gray gray
oxide type refractory dioxide refractory dioxide
oxide density (g/cm³) 4.7 4.7
oxide activity feebly basic feebly basic
chloride boiling point under 100C 86C (GeCl4)
chloride density (g/cm³) 1.9 1.9

• Germanium was isolated in 1882, and provided the
  best confirmation of the theory up to that time,
  due to its contrasting more clearly with its
  neighboring elements than the two previously
  confirmed predictions of Mendeleev do with theirs.

5
Effective Nuclear Charge

• Protons in the nucleus attract the electrons
• Electrons repel each other.
• So inner electrons push the outer electrons
  (shielding), negating much of the pull of the
  nucleus. Thus higher energy levels means less
  lower effective nuclear charge.
• Zeff Z - S

6
Calculating Zeff

• Na11 1s2 2s2 2p6 3s1 Zeff 11 10 1
• Mg12 1s2 2s2 2p6 3s2 Zeff 12 10 2
• Cl17 1s2 2s2 2p6 3s2 3p5 Zeff 17 10 7
• As you can see, the outer electrons of chlorine
  are pulled more by the nucleus than those of the
  sodium or magnesium.

7
Isoelectronic Atoms/Ions

• Iso same
• electronic from electrons
• Isoelectronic particles are those with the same
  of electrons in the same configuration.

8
Size of Atoms
9
Atomic Size on the Periodic Table

• As we compare elements in a period, the Zeff
  increases which means that the valence electrons
  are being pulled harder by the nucleus. So, from
  left to right, the atomic size decreases.

10
Atomic Size on the Periodic Table

• As we compare elements in a family, the main
  difference is the number of shells. From top to
  bottom, the number of shells increases, so the
  atomic size increases.

11
Do Now

• What is the effective nuclear charge of
• An electron in the 3rd energy level of Mo?
• An electron in the 2nd energy level of S?
• An electron in the 4th energy level of Br?
• List these elements in size order P, S, As, Se
• List these particles in size order S, S2-, O

12
Do Now Answers

• What is the effective nuclear charge of
• An electron in the 3rd energy level of Mo?
• Zeff42-1032
• An electron in the 2nd energy level of S?
• Zeff 16 - 2 14

13
Do Now Answers

• An electron in the 4th energy level of Br?
• Zeff35-287
• List these elements in size order P, S, As, Se
• S, P, Se, As
• List these particles in size order S, S2-, O
• O, S, S2-

14
Bond Length

• When a bond forms, two atoms are held next to
  each other by electrical attractions. So the
  distance from nucleus to nucleus is called the
  bond length.
• Bond length is thus the sum of atomic radii.
• For example a C-H bond has a length of 1.14A,
  because C has a radius of .77A and H has a radius
  of .37A. .37A .77A 1.14A.

15
Chart of Atomic Radii
16
What is the bond length of

• C-S?
• S-H?
• N-Cl?
• Na-Cl?

17
What is the bond length of

• C-S 1.79A
• S-H 1.39A
• N-Cl 1.74A
• Na-Cl 2.79A

18
Why is the bond length of NaCl 2.79A?

• NaCl is an ionic compound and thus depends on the
  radii of the ions, not the atoms!
• Na has a radius of .98A and Cl- has a radius of
  1.81A. Thus the sum is 2.79A!

19
Ionic Radii
20
Ionization Energy

• Ionization energy is the minimum energy required
  to remove an electron from the ground state of an
  isolated gaseous atom, or ion.
• Na(g) --gt Na (g) e- IE 496 kJ/mol
• Na (g) --gt Na2 (g) e- IE 4560 kJ/mol
• Why does the first electron come from sodium so
  much easier than the 2nd?

21
Because...

• Na11 1s2 2s2 2p6 3s1
• The first electron comes from the 3rd energy
  level, but the next electron must come from a
  lower energy level, closer to the nucleus, with a
  higher Zeff. Thus it takes a lot more energy to
  get 2 electrons than 1 from a sodium atom.

22
So think about this...

• An element in the 3rd period requires 787 kJ/mol
  to remove its first electron.
• It requires 1575 kJ/mol to remove the 2nd
  electron.
• It requires 3220 kJ/mol to remove the 3rd
  electron.
• It requires 4350 kJ/mol to remove the 4th
  electron.
• It requires 16,100kJ/mol to remove the 5th
  electron.
• What element is this?

23
Chart of Successive Ionizations
The Answer is Silicon
24
Ionization and the Periodic Table

• It is easier to remove a valence electron from a
  bigger element than a smaller one. Why?
• A valence electron in a smaller atom is closer to
  the nucleus, and thus held more tightly by
  electrical attraction.

25
Ionization across a Period

• We now know that the size of the atoms decreases
  as we compare the elements going from left to
  right across a period. This means that more
  energy is required to remove electrons from
  elements on the right (nonmetals) and less for
  elements on the left (metals).
• Ionization energy increases from left to right.
• Could this be why metals give off electrons
  easily?

26
Ionization within a Family

• We also know that each successive member of a
  family is larger because of additional energy
  levels. This means that elements near the top of
  the periodic table require more energy to remove
  an electron than elements near the top.
• Ionization energy decreases from top to bottom.

27
Ionization as a Periodic Function
28
Electron Affinity

• Instead of taking electrons, we could also add
  electrons. One such property of atoms is called
  Electron Affinity.
• Electron affinity is electron affinity is the
  energy released when 1 mole of gaseous atoms each
  acquire an electron to form 1 mole of gaseous 1-
  ions.
• Like Ionization energy, there are successive
  electron affinities.

29
However...

• Electron affinity is not a clear periodic
  property like ionization energy. The reason is
  that energy shifts based on whether subshells or
  orbitals are partially filled or completely
  filled. This makes it hard to come up with a
  good rule.
• Still, it makes sense that a smaller atom can
  attract electrons better than a larger atom. So
  more energy is released when a smaller atom
  captures an electron than a larger atom.

30
Electron Affinity and the Periodic Table

• If smaller atoms release more energy, than
  electron affinity should increase from left to
  right across a period.
• Likewise, electron affinity should decrease from
  top to bottom.

31
Electron Affinity and the Periodic Table
32
Metals
What defines a metal? Weve used words like
luster, ductility, malleability, conductivity.
Why do metals behave this way?
33
Metallic Behavior
Metals tend to be larger atoms. Since it is
easier to remove an electron from a larger atom,
it should make sense then that metals tend to
form cations. Conversely we can say that the
larger an atom is (or the lower its first
ionization is) the more metallic the atom is. So
if we compared O, S, and Se (all nonmetals) we
could say that selenium, being the largest atom,
is the most metallic - even though it is a
nonmetal.
34
Nonmetallic Behavior
Nonmetals tend to be smaller atoms. Since it is
easier to add an electron to a smaller atom, it
should make sense then that nonmetals tend to
form anions. Conversely we can say that the
smaller an atom is (or the higher its first
ionization is) the more nonmetallic the atom
is. So if we compared Li, Na, and K (all metals)
we could say that lithium, being the smallest
atom, is the most nonmetallic - even though it is
a metal.