The Periodic Chart

1
The Periodic Chart

• From then to Now . . .

2
(No Transcript)
3
The History

• 1669-Henning Brand discovered Phosphorus
• 1680-Robert Boyle rediscovered Phosphorus
• 1789-Lavosier wrote the 1st chemistry text
• 1809-There were 47 known elements
• 1862-Beguyer deChancourtois noticed periodicity
• 1863-Newlands classified 56 elements into 11
  groups, octaves
• 1869-Mendeleev created a table and was able to
  predict the existence of 2 new elements
• 1869-Meyer also created a table, but did not get
  the credit for it
• 1900-Moseley developed the Periodic Law
• 1944-Seaborg proposed the Actinide series

4
History of the periodic table

• In the 1700s only 30 elements were identified
• Dobereiner in the 1800s noticed certain elements
  could be grouped into sets of 3 called triads
• Dobereiner--triads

5
TRIAD PROPERTIES

• Properties similar Group 1 are soft metals
• Reactiveness similar Group 1 are very reactive
  with water
• Middle element value is average of one above and
  one below
• Triad3 elements with similar properties
• one value is an average of the other 2

6
PROPERTIES OF TRIADS
Li (3) Atomic number Ca 40.1 amu Atomic mass Cl 1.56 g/cm3 Density
Na (11) Atomic number Sr 87.6 amu Atomic mass Br 3.12 g/cm3 Density
K (19) Atomic Number Ba 137.0 amu Atomic mass I 4.95 g/cm3 Density
7
Newlands mid 1800s

• Now 49 elements
• Noticed that when arranged by increasing mass,
  every 8th element had similar properties
• Called law of octaves
• Newlands -- octaves

8
MENDELEEV VS MEYER

• Both made discoveries at the same time but
  Mendeleev was the first to publish them
• Wrote names and properties on cards and arranged
  them in various ways
• In increasing mass
• In repetitive properties
• Both couldnt be done at the same time
• Decided putting them in order of repetitive
  properties was more important

9
DISCREPANCIES

• In order to put the elements in similar groups
  according to properties, some of the masses were
  out of order
• Thought that the atomic masses were wrong

10
MENDELEEVS PERIODIC TABLE

• When he put elements in order according to their
  properties without regard to their masses, some
  elements seemed to be missing
• He predicted the existence of these missing
  elements and when discovered, they fit perfectly
  into his pattern
• But Mendeleev was not entirely correct
• The atomic masses, when recalibrated, were
  not incorrect.
• This left some atomic masses out of order on his
  periodic table

11
Mendeleevs notes
12
Mendeleevs 1869 Periodic Table
13
LATE 1800S MOSELYS PERIODIC TABLE

• Developed the idea of atomic s
• Assigned one to each element based on the of
  protons in their nucleus
• Arranged elements according to the number of
  protons instead of mass
• Now, elements are in a numerical repetitive order
  as well as grouped according to their properties
• Since masses arent figured into arranging the
  periodic table, its ok for them to be out of
  order

14
The Periodic Law

• The periodic properties of the elements are
  functions of their atomic number.
• In other words, the elements are arranged on the
  basis of their ground state electron configuration

15
Periodic Table of 1944
16
The Modern Periodic Table
17
Vertical Columns

• The vertical columns are arranged in groups or
  families.
• They are numbered from left to right
• Elements in a group have the same electron
  structure in their outer subshell (valence
  electrons)

18
Electron Review

• An electron shell, also known as a main energy
  level, is a group of atomic orbitals with the
  same value of the principal quantum number n.
• Electron shells are made up of one or more
  subshell, which have orbitals with the same
  angular momentum quantum number l. (1 of s, 3 of
  p, 5 of d and 7 of f orbitals)

19

• States with the same value of n are related, and
  said to lie within the same electron shell.
• Example 1s22s22p6
• 1s2 and 2s22p6 are in the same electron shell
• States with the same value of n and also l are
  said to lie within the same electron subshell.
• Example 1s22s22p6
• 1s2 are in the same electron subshell
• 2s2 are in the same electron subshell
• 2p6 are in the same electron subshell

20

• Electron shells make up the electron
  configuration.
• It can be shown that the number of electrons that
  can reside in a shell is equal to 2n2.
• Shells and subshells are defined by the quantum
  numbers.
• In large atoms, shells above the second shell
  overlap (Aufbau principle)

21
Valence Shell

• The valence shell is the outermost shell of an
  atom, which contains the electrons most likely to
  participate in a chemical reaction with other
  atoms or to determine chemical properties.
• Electrons in the valence shell are referred to as
  valence electrons.

22
Lets see ...
Group 18 Ne 1s22s22p6 Ar 1s22s22p63s23p6 Kr
Ar4s23d104p6 Xe Kr5s24d105p6 Rn
Xe6s24f145d106p6
23
Further Breakdown
s-orbital elements d-orbital elements p-orbital
elements
f-orbital elements
24
Horizontal Rows

• The horizontal rows are the periods.
• The periods are numbered from the top down.
• Elements in the same period have the same
  principal energy level

25
Lets see ...
Period 2 (Period n) Li 1s22s1 Be 1s22s2
B 1s22s22p1 C 1s22s22p2
26
Group Names

• Groups 1-2 and 13-18 (except Hydrogen) are the
  main group elements (also known as the
  representative elements).
• Groups 3-12 are the transition metals

27
Specific Group Names

• Group 1 alkali metals
• Group 2 alkaline earth metals
• Group 11 coinage metals (not IUPAC approved)
• Group 15 pnictogens (not IUPAC approved)
• Group 16 chalcogens
• Group 17 halogens
• Group 18 noble gases

28
Period Identifications

• The elements in the 1st f-period are the
  Lanthanide series.
• The elements in the 2nd f-period are the Actinide
  series

29
Group 1 Alkali Metals

• Hydrogen is NOT included in Group 1
• Metals that react with water to make an alkaline
  solution (basic)
• Highly reactive, soft (less than 1 on the Mohs
  scale), and conductive

30
Group 1 Electrons

• Not found in their elemental form but in
  compounds
• example NaCl, KOH
• There is only 1 valence electron. (ns1)
• If the one electron is lost, it will be stable

31
Mohs Hardness Scale

• The scale used to describe the hardness of a
  material is the Mohs Hardness Scale
• The scale is from 0-10 (softest to hardest)
• example Talc is 1 on the Mohs scale and the
  Diamond is 10

32
Group 2 Alkaline Earth Metals

• The alkaline earth metals are silvery colored,
  soft, low-density metals, which react readily
  with halogens to form ionic salts, and with
  water, to form strongly alkaline hydroxides.
• Highly reactive, but not as reactive as alkali
  metals, usually found as compounds not in
  elemental form

33
Alkaline Earth Electrons

• There are 2 valence electrons. (ns2)
• It takes more energy to lose 2 electrons than it
  does to lose only one (like the alkali metals)

34
Valence Electrons of Groups 13-18

• Group 13 ns2np1
• Group 14 ns2np2
• Group 15 ns2np3
• Group 16 ns2np4
• Group 17 (halogens) ns2np5
• Group 18 (noble gases) ns2np6

35
Group 17 The Halogens

• Halogens are highly reactive non-metals.
• Only 7 valence electrons (just one short of a
  full and stable valence shell) so they want to
  gain an electron
• Reactive with most metals to form salts

36
Group 18 Noble Gases

• Have a full set of electrons (n2p6)
• Low chemical reactivity and so they are very
  stable

37
Hydrogen

• Hydrogen is in a class by itself because it is
  the most common element in the Universe!
• Hydrogen only has one proton and one electron and
  can react with almost anything

38
Transition Metals

• Groups 3-12 (d-block)
• Do NOT have identical electron configurations in
  the outer shell. Why?
• The Lanthanide and Actinide series are contained
  within the d-block and have f-orbitals

39
Lanthanide Actinide

• Lanthanide are the rare earth series from atomic
  58 to 71
• shiny metals with similar reactivity to alkaline
• Actinide are from atomic 89 to 103
• nuclei are unstable, radioactive
• As you move to the right, electrons are filled in
  the f-orbital

40
Metallic Character

• Approximately 2/3s of the elements are metals.
• See periodic chart
• Metals have unique properties
• luster mirror like shine that reflects light
• conductivity ability to conduct heat or
  electricity
• malleable ability to be rolled or hammered
• ductile ability to be drawn into wire

41
Alloys

• Metals that are mixed with other metals to form a
  stable compound are called alloys
• example Brass is Copper and Zinc
• example Steel is Iron, Tin, Nickel, Lead, etc.

42
Nonmetals

• Poor conductors of heat and electricity
• Not malleable
• Many are gasses
• One is liquid Br
• Some are solids (brittle and dull)
• More electrons in outer level
• Form negatively charged ions

43
METALLOIDS

• Metalloids have properties of both
• metals and nonmetals
• On the stairstep exclude Aluminum and Polonium