Periodic Table

1
Periodic Table

• 1869
• Dmitri Mendeleev in Russia
• Lothar Meyer of Germany
• Published nearly identical schemes

2
Henry Moseley

• 1913
• Developed the concept of atomic numbers

3
Groups

• Vertical Columns
• Families
• Elements have similar properties
• Same number of valence electrons (the number of
  electrons in outer most shell)

4
Periods

• Horizontal rows
• Period number corresponds to number of energy
  levels

5
Period   GROUPIA   GROUPIIA Transition Metals Transition Metals Transition Metals Transition Metals Transition Metals Transition Metals Transition Metals Transition Metals Transition Metals Transition Metals GROUPIIIA GROUPIVA GROUPVA GROUPVIA   GROUPVIIA   GROUPVIII (0)
1 2He-
  2 3Li1.0 4Be1.5 5B2.0 6C2.5 7N3.0 8O3.5 9F4.0 10Ne-
3 11Na0.9 12Mg1.2 13Al1.5 14Si1.8 15P2.1 16S2.5 17Cl3.0 18Ar-
4 19K0.8 20Ca1.0 21Sc1.3 22Ti1.5 23V1.6 24Cr1.6 25Mn1.5 26Fe1.8 27Co1.9 28Ni1.9 29Cu1.9 30Zn1.6 31Ga1.6 32Ge1.8 33As2.0 34Se2.4 35Br2.8 36Kr-
5 37Rb0.8 38Sr1.0 39Y1.2 40Zr1.4 41Nb1.6 42Mo1.8 43Tc1.9 44Ru2.2 45Rh2.2 46Pd2.2 47Ag1.9 48Cd1.7 49In1.7 50Sn1.8 51Sb1.9 52Te2.1 53I2.5 54Xe-
6 55Cs0.7 56Ba0.9 57La1.1 72Hf1.3 73Ta1.5 74W1.7 75Re1.9 76Os2.2 77Ir2.2 78Pt2.2 79Au2.4 80Hg1.9 81Tl1.8 82Pb1.9 83Bi1.9 84Po2.0 85At2.2 86Rn-
7 87Fr0.7 88Ra0.9 89Ac1.1 104Rf- 105Ha-
6
Periodic Table

• Most of the elements are metals
• Gases H,N,O,F,Cl and Group 18
• Liquids Hg and Br

7
Metals

• Left of zig zag line
• Low ionization energy and electronegativity
• Lose electrons to form positive ions
• Metallic luster when polished
• Good conductor of heat and electricity
• Malleable (made into sheets)
• Ductile (made into wire)
• Most are solid at room Temperature (Hg is a
  liquid)

8
Lower Left of Table

• Most Metallic Character
• Most reactive metal
• YouTube - Brainiac Alkali Metals

9
Nonmetals

• Right of zig zag line (except group 18 Noble
  gases)
• Gain electrons to form form negative ions
• High ionization energy (attraction for electrons)
  and high electronegativity
• Lack luster and form brittle solids
• Poor conductors of heat and electricity
• Exist as gases (mostly) solids (molecular or
  network) bromine is a volatile liquid at room
  temperature.

10
Metalloids

• Touch zig zag
• Known as semi-metals
• At the border of metals and nonmetals
• Properties are similar to metals and nonmetals
• Boron, silicon, germanium, arsenic, antimony and
  tellurium

11
Metalloids
12
Group 1 Elements

• Alkali Metals
• One valence electron
• React with acids to product hydrogen gas
• Form strong bases

13
Group 2 Metals

• Alkali earth metals
• 2 valence electrons
• Form weak bases

14
Metals

• Reactivity increases as we go down a group
• Reactivity decreases as we go across a period
  (left to right)
• Ex. Sodium is more reactive than Magnesium
  (period 3)
• Why? Sodium wants to give the one electron away
  faster.
• Group 1 larger atomic size
• Group 2 greater nuclear charge

15
Group 17

• Halogens
• 7 valence electrons
• Diatomic molecules
• High electronegativity
• Very reactive, occurs in nature only as compounds
• All three phases of matter exist.

16
Phases

• Fluorine and chlorine are gases
• Br is a liquid
• Iodine and astatine are solids

17
Group 18

• Noble Gases
• Inert Gases-old name, they were thought to be
  unreactive.
• Chemically unreactive
• Completely filled valence shells
• Monoatomic

18
Groups 3-11

• Transition elements
• Positive oxidation states
• Formed colored compounds
• d sublevels lose electrons from an inner energy
  level
• Forms ions of more than one stable charge
• Ex. Fe2 and Fe3

19
Trends

• Bottom left most reactive metal
• Top right (excluding group 18) most reactive
  nonmetal
• Metals reactivity increases as you go down a
  group
• Nonmetals reactivity increases as you go up a
  group

20
Ionization Energy

• Energy needed to remove an electron from an atom
• The greater the attraction for electrons the more
  energy is needed to remove the electron
• The smaller the ionization energy, the easier it
  is to remove an electron

21
Ionization Energy

• Li 520
• B 801
• O 1314
• Ne 2081
• Trend Ionization Energy increases across a
  period.

22
Ionization Energy

• C 1086
• Si 787
• Ge 762
• Pb 716
• Trend Ionization Energy decreases down a group.

23
Electronegativity

• Is the attraction for electrons. The larger the
  electronegativity the more the atom attracts
  electrons.
• 3 Factors
• Nuclear charge - of protons, increase charge
  greater attraction
• Principal energy level The higher the principal
  energy level of the outermost electrons, the
  greater the distance from the nucleus, the weaker
  is the pull
• Electron Cloud Effect shields the outermost
  electrons from the nucleus. The inner electrons
  repel outermost electron

24
Electronegativity

• Li 1.0
• B 2.0
• O 3.4
• Ne -----
• Why doesnt Ne have an Electronegativity?
• It is not reactive. Does not react with any
  element.
• Trend Electronegativity increases across a period.

25
Electronegativity

• C 2.6
• Si 1.9
• Ge 2.0
• Pb 1.8

• F 4
• Cl 3.2
• Br 3.0
• I 2.7

26
Electronegativity

• Trend
• Electronegativity decreases down a group.

27
Atomic Radius

• Is half the distance between adjacent nuclei (the
  distance from the nucleus to the outer most
  valence electrons).
• Size of the atom
• Related to the attraction of the nucleus for its
  electrons
• Decreases as the atomic number increase as a
  result of the force of attraction between the
  positive nucleus and the negative electrons.

28
Atomic Radius

• Li 130
• B 84
• O 64
• Ne 62
• Trend Atomic Radius decreases across a period.

29
Atomic Radius

• C 75
• Si 114
• Ge 120
• Pb 145
• Trend Atomic Radius increases down a group.

30
Ionic Radius

• Metals lose electrons, ionic radius becomes
  smaller
• Nonmetals gain electrons, the ionic radius
  becomes larger

31
What do all of these have in common?
32
They are all Carbon.
33
What do all of these have in common?
Graphite
Carbon Fiber Clothes
Artists Charcoal
Carbon Fiber Custom Wheel
Charcoal
Diamond
34
Allotropes

• Arent we all the same deep down?

35
What is an allotrope?

• A structurally differentiated form of an element
  that exhibits allotropy.

36
What is allotropy?

• A property of certain elements, as carbon,
  sulfur, and phosphorus, of existing in two or
  more distinct forms.

37
A.K.A.

• When an element exists in pure forms that differ
  in the way the atoms are arranged.

38
Buckminsterfullerene
39
(No Transcript)
40
Other Allotropes

• Sulfur
• S4, S5, S8
• Oxygen
• O2 and O3
• Phosphorous
• P4 and P2

41
(No Transcript)
42
(No Transcript)
43
(No Transcript)
44
(No Transcript)
45
(No Transcript)