Unit 4 The Periodic Table of the Elements I' History of the Periodic Table Chapter 6 p155 to 163 and

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Unit 4 The Periodic Table of the ElementsI.
History of the Periodic Table Chapter 6 p155 to
163 and 166 to 178

• A. Johann Dobereiner 1829
• Some elements had similar properties
• Grouped elements in threes, called Triads
• Ex. Li, Na, K Cl, Br, I
• B. John Newlands -1864
• Organized elements by increasing mass
• Noticed every 8th element had similar properties
• Li Be B C N O F
  Na Mg
• As new elements discovered, this pattern failed.

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C. Dimitri Mendeleev 1869 Created the first
Periodic Table of elements based on two rules

• Elements arranged from lightest to heaviest.

2. Columns contain elements with similar
properties
His genius was shown in two ways

• He placed similar elements in the same column
  even if they ended up out of weight order. ex.
  Iodine

2. He left blank spaces in his table for
elements not yet discovered and was able to
accurately predict their properties! ex.
ekasilicon Germanium
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D. Henry Moseley - 1914

• Used X-ray crystallography to discover atomic
  number the number of protons.
• Arranged elements by increasing atomic number
  instead of atomic mass.
• When lined up this way, columns contained similar
  elements with no exceptions.
• Created the MODERN PERIODIC LAW
• Properties of Elements are periodic functions
  of their atomic numbers

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• E. Organization of the Modern Periodic Table
• Groups the columns on the table.
• Include elements with similar properties.
  Because of this, groups are also called Families.
• All elements in the group have the same number of
  valence electrons. (This is why they are
  similar).
• Remember Number of valence electrons column
  number without teen.

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• 2. Periods the rows on the periodic table.
• 1st period is H and He, 2nd is Li, Be,, Ne.
• The Period number tells you how many principle
  energy levels are used.
• ex. K is in the 4th row and has 4 energy
  levels
• The bottom two rows actually fit into rows 6 and
  7 (they are called the Lanthanides and Actinides
  respectively). They are cut out to make the
  table shorter!

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II. Metals, Nonmetals and Metalloids

• Metals elements to the left of the stair step
  line on the Periodic Table.

(except Hydrogen!)
Properties of Metals

• All metals have loose electrons
• 2. Good conductors of heat and
• electricity.
• 3. Have metallic luster (shiny).

Metals
4. All are solids at room temperature except
mercury (a liquid).
5. Are ductile (can be drawn into a wire) and
malleable (can be pounded into sheets).
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B. Nonmetals All elements to the right of the
stair step line and hydrogen.
(except column 18)
Properties of Nonmetals
1. Hold on to their valence electrons strongly
2. All are dull and brittle in solid form.
3. Do not conduct electricity or heat very well.
4. At room temperature, most are solids but some
are gases (H, N, O, F, Cl) and one is a liquid
(Bromine).
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C. Metallic vs. Nonmetallic Character

• Metallic Character how much the element acts
  like a metal(Shininess, conductivity etc.)
• Metallic Character increases as you go down the
  table and to the left Francium has the most
  metallic character.
• Nonmetallic Character increases as you go up and
  to the right - Fluorine has the most nonmetallic
  character.
  Nonmetallic
• 
  Metallic

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D. Metalloids (Semimetals)

• Metals that have some properties of nonmetals
• Nonmetals that have some properties of metals.
• Most of the elements on the stair step line are
  metalloids

B Boron
Si Silicon
Germanium Ge As Arsenic
Antimony Sb Te Tellurium
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E. Reaction of Metals with Nonmetalstransfer of
electrons

• Filled energy levels are most stable (lowest
  energy)
• Most elements want 8 valence electrons (the octet
  rule) Hydrogen and Helium only need 2.
• Noble gases (column 18) have filled last energy
  levels and do not react.
• Metals lose all valence electrons and revert to
  an inner filled energy level as the last level
  become positive ions (cations).
• Nonmetals gain enough electrons to make 8 valence
  electrons become negative ions (anions).

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• When metals and nonmetals react with each other,
  the nonmetal takes electrons from the metal. The
  ions formed attract each other and an ionic bond
  is formed.
• Ex. Na and Cl

Transfer of electron
Na Cl

Ions form that attract each other.
Na
Cl

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IV. Characteristics of Some Groups

• Group 1 Alkali Metals Base forming
• Have 1 valence electron and 1 charge.
• Most reactive metals Never found in pure form
  (uncombined) in nature.
• Hydrogen (H2) not included in this group.
• Group 2 Alkaline Earth Metals common on earth
• Have 2 valence electrons and 2 charge.
• Very reactive metals Never found uncombined in
  nature.

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C. Group 17 Halogens stinky elements

• Most reactive nonmetals
• All exist as diatomic molecules in pure form
  Fluorine (F2) and Chlorine (Cl2) are gases,
  Bromine (Br2) is the only nonmetal liquid, and
  Iodine (I2) and Astatine (At2) are solids
• Only column with All Three Phases
• Only column with More Than One Diatomic

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D. Group 18 Noble Gases

• All have full valence energy level, usually an
  octet (8), Helium has 2.
• Do not react with other elements.
• Exist as Monoatomic gases at room temp.
• Kr and Xe can be forced to react (with O2 and F2)

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E. Transition Elements Groups 3-11

• All have 2 unfilled levels (remember overlap)
• Allows electrons from two different levels to
  react.
• Therefore, most have several possible charges
  (oxidation states) and can make more than one
  possible compound with nonmetals.
• Transition metals make colored compounds.
• ex. Iron Oxide FeO is red and Fe2O3 is
  orange

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IV. Property Trends on the Periodic Table
A. Atomic Radius Atomic Size half the
distance between nuclei of two atoms of an
element bonded together.
radius distance
2
distance

• Atomic radius increases as you go down a column
  on the periodic table. Why?

more energy levels!
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1
3
H 1
Li 2-1
Na 2-8-1
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• 2. Atomic radius (and atom size) decreases as you
  go left to right across a period. Why?

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5
3
6
Li 2-1
Be 2-2
B 2-3
C 2-4
While the number of energy levels remains the
same, the nuclear charge ( the number of protons)
increases, pulling the electrons closer to the
nucleus, and the atoms get smaller.
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• B. Ionic radius radius of an atom after it has
  gained or lost electron(s) to form an ion.

1. Metals lose electrons to form positive
cations. How does the radius of the cation
compare to the atom.
The cation is smaller because it has fewer
electrons and one less energy level!
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ex. Li
Li
2. Nonmetals gain electrons to form negative
anions. How does the radius of the anion compare
to the atom?
The anion is larger because it has the same
number of protons but more electrons, which are
held more loosely!
ex. O
O2-
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C. Ionization Energy (I.E.) the energy needed
to remove the outermost electron from an atomto
make a positive ion (cation).

• See table S ex. Look up values for Na and K
• Na(g) 496 kJ/mole ? Na 1e
  (electron)
• K(g) 419 kJ/mole ? K 1e
• Potassium (K) is more reactive than Sodium (Na),
  it has looser electrons as seen by its lower I.E.
• Lower I.E. ? More reactive metal.
• Notice (Act 3-9) as you go down a group, I.E.
  decreases (atomic radius increases). Why?
• larger atomic radius ? valence es further
  from nucleus ?es more loosely held? Lower I.E.

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• Notice as you go from left to right across a
  period, I.E. increases. Why?
• smaller atom ? valence es closer to nucleus ?
  electrons more strongly held ? higher I.E.
• D. Electronegativity(Eneg) a measure of an
  atoms ability to attract electrons.
• It is a relative scale that goes from 0 to 4
• 0 no ability to attract electrons
• 4 greatest ability to attract electrons
• Noble gases have no Eneg
• Lowest Eneg most reactive metals Fr Cs 0.7
• Highest Eneg most reactive nonmetals F 4.0

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• Why does Eneg increase as you go toward fluorine
  (either up a column or left to right)?
• Atom size (radius) decreases toward Fluorine.
  The nucleus of an atom attracts electrons. The
  smaller an atom the closer its nucleus can get to
  another atoms valence electrons and attract them!

Fluorine is smaller and can get closer to
Potassiums valence electron
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Also, Fluorine has fewer electrons overall, less
electron shielding.