A German scientist named Johann Dobereiner put forward his law of triads in 1817.

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A German scientist named Johann Dobereiner put
forward his law of triads in 1817.
Each of Dobereiner's triads was a group of three
elements. The appearance and reactions of the
elements in a triad were similar to each other.
At this time, scientists had begun to find out
the relative atomic masses of the elements.
Dobereiner discovered that the relative atomic
mass of the middle element in each triad was
close to the average of the relative atomic
masses of the other two elements.

This suggested that atomic mass might be
important to arranging the elements.
This gave other scientists a clue that relative
atomic masses were important when arranging the
elements
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With this idea in mind nearly 50 years later, in
1864, an Englishman, John Newlands, arranged the
known elements into seven rows.
Newlands' Octave Arrangementfrom Chemical News
1866
No. No. No. No. No. No. No. No.
H 1Li 2G 3Bo 4C 5N 6O 7 F 8Na 9Mg 10Al 11Si 12P 13S 14 Cl 15K 16Ca 17Cr 19Ti 18Mn 20Fe 21 Co Ni 22Cu 23Zn 25Y 24In 26As 27Se 28 Br 29Rb 30Sr 31Ce La 33Zr 32Di Mo 34Ro Ru 35 Pd 36Ag 37Cd 38U 40Sn 39Sb 41Te 43 I 42Cs 44Ba V 45Ta 46W 47Nb 48Au 49 Pt Ir 50Os 51Hg 52Tl 53Pb 54Bi 55Th 56
The elements were placed in order of increasing
atomic mass and arranged so elements with similar
chemical properties are in the same group.
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Since the properties were repeated every eighth
element, Newlands referred to his arrangement as
the
Law of Octaves.
Unfortunately for
Newlands, there were a number of problems with
his arrangement and it really only worked up
through calcium.
About five years later, Dmitri Mendeleev a
Russian chemist who was unaware of the work of
Newlands,
designed his own table.
After about 18 months of gathering
information and arranging element
cards, it was finished.
Like Newlands, Mendeleev arranged the 63 known
elements in order of increasing atomic mass,
having elements with similar chemical properties
in the same group.
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Unlike Newlands his table had groups of varying
lengths.
He also left gaps in the table for elements he
believed had not yet been discovered.
He
even made predictions about the properties of
some of these elements.
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Although his successful predictions allowed many
to accept his periodic idea,
there were
still anomalies that Mendeleev could not explain.
One of these is the order of iodine and tellurium.
The problem was finally solved in 1913 by a
25-year-old English physicist named Henry Moseley.
Moseley showed that the ordering of X-ray
spectral lines was dependent upon the ordering of
nuclear charge,
that is, in order of the atomic number.
When the elements were placed in order of
increasing atomic number,
the anomalies in Mendeleevs table were
eliminated.
Moseleys work gave rise to the modern periodic
law
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The properties of the elements are a periodic
function of their increasing atomic numbers.
Tragically for the development of science,
Moseley was killed in battle during World War I,
only two years later.
Moseleys revised periodic table looked something
like this
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The Modern Periodic Table

• seven horizontal rows called periods
• 18 vertical columns called groups or families
• groups 1 and 2 and groups 13-18 are called
  representative elements
• groups 3-12 are the transition metals

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Modern Periodic Table cont.

• elements in any group have similar physical and
  chemical properties
• properties of elements in periods change from
  group to group
• symbol placed in a square
• atomic number above the symbol
• atomic mass below the symbol

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Metallic Character

• Metals
• malleable ductile
• shiny, lustrous
• conduct heat and electricity
• lose electrons in reactions

• Nonmetals
• brittle in solid state
• dull
• electrical and thermal insulators
• gain electrons in reactions

• Metalloids
• Also known as semi-metals
• Show some metal and some nonmetal properties

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Metallic Character

• Metals are found on the left of the table,
  nonmetals on the right, and metalloids in between
• Most metallic element always to the left of the
  Period, least metallic to the right, and 1 or 2
  metalloids are in the middle
• Most metallic element always at the bottom of a
  column, least metallic on the top, and 1 or 2
  metalloids are in the middle of columns 4A, 5A,
  and 6A

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Other Important Groups to Know

• Group IA ? alkali metals
• Group IIA ? alkaline earth metals
• Group VIIIA ? noble gases
• Group VIIA ? halogens salt formers
• Group VIA ? chalcogens
• Group VA ? Nitrogen group
• Group IVA ? IVA group
• Group IIIA ? IIIA group

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Other Groups

• s p block filling ? representative elements
• d block filling ? transition metals
• f block filling ? inner transition metals
• 4f ? lanthanides
• 5f ? actinides
• f elements that are naturally occurring ? rare
  earth elements

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What are Periodic Trends

• trends are general patterns or tendencies
• they are general not definite there are
  exceptions
• when looking at trends we look for increases
  decreases
• across ? periodic
• down ? group

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Effects on the Trends

• Nuclear Charge
• the pull of the nucleus
• proportional to the number of protons in an atom
• the greater the number of protons, the stronger
  the nuclear charge (pull)
• this has its greatest effect across a period

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Effects on Trends Cont.

• Shielding
• - the electron protection from the nuclear
  pull
• - shield an energy level of electrons
• - we are not concerned with single electrons,
  only energy levels of electrons
• - these electrons reduce the nuclear pull
• - affects group trends

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Effects on Trends Cont.

• Stability
• - where electron arrangement is compared to
  stable octet (or other special stabilities)
• - determines if atom gains or loses electrons
• - can be used to explain anomalies in trends

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Trend in Atomic Size

• Decreases across period
• left to right because of the nuclear pull
• adding electrons to same valence shell
• valence shell held closer because more protons in
  nucleus
• Increases down column
• valence shell farther from nucleus because of
  increased shielding
• Illustration

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Trend in Ionization Energy

• Minimum energy needed to remove a valence
  electron from an atom
• 1 mole of electrons in the gaseous state (kJ/mol)
• The lower the ionization energy, the easier it is
  to remove the electron
• metals have low ionization energies
• Ionization Energy decreases down the group
• valence electron farther from nucleus
• Ionization Energy increases across the period
• left to right
• harder to remove an electron from the atom
  because of the increased nuclear pull
• Exceptions Group 3 less than Group 2, Group 6
  (chalcogens) less than Group 5

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Ionization Energy Cont.

• Li energy ? Li e-
• 1st ionization 520 kJ/mol
• Li energy ? Li2 e-
• 2nd ionization 7297 kJ/mol
• Li2 energy ? Li3 e-
• 3rd ionization 11,810 kJ/mol
• Notice, each successive ionization energy is
  greater than the preceding one there is a
  greater pull between the nucleus and the
  electron and thus more energy is needed to break
  the attraction.
• Examining ionization energies can help you
• predict what ions the element will form.
• easy to remove an electron from Group IA, but
  difficult to remove a second electron. So group
  IA metals form ions with a 1 charge.

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Ionization Energies for aluminum (kj/mol)
1st - 578
2nd - 1817
3rd - 2745
4th - 11580
5th - 14840
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Electron Affinity

• atoms tendency to attract (gain) an electron
• it is the energy change that accompanies the
• addition of an electron to a gaseous atom
• Basically the opposite of ionization energy
• Across a Period
• electron affinity increases because of increased
  pull
• Down a Group
• electron affinity decreases because the electrons
  are shielded from the pull of the nucleus
• Exceptions 2A Nitrogen Group Noble Gases

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Ionic Size

• cations lose electrons (positively charged)
• anions gain electrons (negatively charged)
• elements gain or lose e- to become stable being
  like noble gases (filled outer sublevel)
• IA - 1 VA - -3
• IIA - 2 VIA - -2
• IIIA - 3 VIIA - -1
• IVA share VIIIA 0, stable
• Illustration

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Ionic Size Cont.

• Across a Period
• cations decrease (I-III) because of greater pull
  on electrons (protons pull on fewer electrons
  more )
• anions decrease (V-VII) because of decrease in
  electron repulsion (protons pull on fewer
  electrons i. e. less negative)
• Down a Group
• both cations and anions increase size
• GOOD RULE OF THUMB
• anions are always larger than their neutral atom
• cations are always smaller than neutral atom

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Electronegativity

• the ability of an atom to attract electrons when
  the atom is in a compound
• electron tug-of-war
• similar to electron affinity, but not the same
• Across a Period
• increases because of increased pull
• Down a Group
• decreases because of shielding
• Fluorine most electronegative element

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Reactivity

• Reactivity of metals increases to the left on the
  Period and down in the column
• follows ease of losing an electron
• Reactivity of nonmetals (excluding the noble
  gases) increases to the right on the Period and
  up in the column
• Reactivity Video

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Practice

• Which element has a greater ionization energy
  Mg or Ba
• Which element has a greater atomic radius N or
  F
• Which element has a greater electron affinity S
  or Pb