Download Periodic Table Development

Periodic Table Development
Late 1700’s – 30 elements discovered
1820 – Dobereiner’s Triads
 Grouped elements into sets of 3 (triads)
(Li, Na, K) ( Ca, Sr, Ba ) ( Cl, Br, I )
Properties of the middle element are averages of the 1st and 3rd
1865 – Newland’s Octaves
 Organized elements into repeating groups of 8
Law of Octaves – elements arranged by increasing atomic mass;
properties of the 8th element are similar to the 1st
1869 – Mendeleev
 Arranged elements by increasing atomic mass
 Observed periodic (repeating) element properties
 Produced 1st Periodic Table
 Elements in the same column had similar properties
 Predicted properties of undiscovered elements
1913 – Moseley
 Arranged elements by atomic number
Periodic Law –
Elements arranged in order of increasing atomic number show a periodic
pattern in their physical and chemical properties
Reading the Periodic Table
Information in each square:
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Element name
Element symbol
Atomic # = ( # protons or # electrons)
Atomic Weight = (weighted average of isotope masses)
Electron configuration
Group = elements in vertical column with similar properties. (Families)
Period = Horizontal row of elements
18 labeled groups and 7 periods
Labeling and Naming Groups
Group number
Family Name
Group 1A =
Group 2A =
Group 7A =
Group 8A =
Alkali metals
Alkaline Earth metals
Halogens
Noble gases
Metals, Nonmetals, Semimetals
Metal Properties
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Metallic Luster (shine)
Good conductor of heat and electricity
Solids at Room temperature ( except Mercury)
Malleable ( hammered into thin sheets without shattering)
Ductile (drawn into fine wires)
Examples: Cu, Ag, Al, Au, Zn
Most elements are metals
Nonmetals
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No metallic luster (dull)
Poor conductors of electricity and heat
Neither malleable nor ductile (brittle)
Gases and solids at room temperature
Bromine is a liquid at room temperature
Located on far right side of the Periodic Table
C, N, O ,P, S, F, Cl, Br, I. He, Ne, Ar, Kr, Xe
Semimetals or Metalloids
Have properties of both metals and nonmetals
Boron, Silicon, Germanium, Arsenic, Antimony
Why do elements in a group have similar properties?
Similar electron arrangement
Valence electrons
Electrons in the outermost energy level of an atom are
responsible for an atom’s chemical behavior
Elements in the same group have valence electrons in similar electron
configurations
Group 1A (alkali metals)
all have one valence electron in the S orbital
Abbreviated electron configurations
 Focus on valence electrons
Inner core electrons represented by the symbol of the nearest noble gas with a
lower atomic #
Group 1A elements:
H = 1S1
Li = [He] 2S1
Na = [Ne] 3S1
K = [Ar] 4S1
Rb = [Kr] 5S1
Cs = [Xe] 6S1
[He] = 1S2
 Each Group 1A element has a single valence electron in the s orbital
 Principal quantum # of the s orbital = Element’s row or period
S, P, d, f – block elements
Periodic Table is divided into 4 blocks
s – block
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Group 1A and 2 A
Alkali metals & Alkaline Earth metals
Valence electrons in S orbitals only
Group 1A, each ecn ends in S1
Group 2A each ecn ends in S2
p – block
 Group 3A to Group 8A of any period
 Valence electrons in P1 to P6 orbitals
 P sublevel can hold 6 electrons
d- block
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Middle of Periodic Table
d sublevel can hold 10 electrons
3d orbitals start with Sc (atomic # 21)
d-block elements are called transition metals
f – block
 electrons start to be located in f orbitals
 f orbitals can hold 14 electrons
 start filling 4f orbitals on 6 period with La (atomic # 57)
f – block elements called inner transition metals
Periodic Table shape is due to the way electrons fill s,p,d, f orbitals of different
energy levels
Periodic Trends
Systematic changes of element’s properties throughout the periodic table
Properties determined by an atom’s electron configuration
Periodic Trends include:
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Atomic radius
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Ion Size
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Ionization Energy
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Electron Affinity
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Electronegativity
Atomic radius
Distance from center of an atom’s nucleus to its valance electrons
Atomic radii increase moving down a group
Why?
There are more electrons down a group,
Energy levels holding those electrons are farther away from the nucleus
With increasing distance, there is less attractive force exerted by the nucleus on
the electrons.
Therefore, atomic radius increases
Atomic radii decrease moving across a row
from left to right
Why?
Across a period, the increasing numbers of protons exert a stronger pull of the
electrons.
Valence electrons are attracted to the nucleus
This attraction shrinks the electron orbital to reduce the atomic radius
Ion Size
As an atom loses electron(s) to form a positive ion, it becomes smaller
Li atom’s radius is 0.152 nm
Li +1 ion’s radius is 0.060 nm
Loss of electrons vacates the largest orbital
Atom gains an electron to form a negative ion, it becomes larger
Fluorine atom’s radius = 0.064 nm
Fluorine ion’s radius = 0.136 nm
Increasing the # electrons increases repulsive force to spread electron cloud
Periodic trend, elements in the same group form ions of the same charge
Ionization Energy ( I.E.)
Energy needed to remove an electron from an atom
Li (vapor) + Energy ---------- Li +1 (vapor) + electron
I.E. represents how strongly an atom holds onto its valence electrons
Periodic Trend:
I.E. decreases down a group
I.E. Increases across a period ( Left to Right)
Both atomic radii and I.E. depends on how strongly an atom’s electrons are
attracted to the nucleus
Electron Affinity ( E.A.)
Energy released when an atom gains an extra electron
Ne (gas) + electron --------- Ne –1 (gas) + Energy
E.A. represents the atom’s attraction for an extra electron
An atom’s E.A. is related to the # of electrons it needs to fill its outer energy
level
Nonmetals E.A. > Metal E.A.
Electronegativity
Ability of an atom to attract an electron in a chemical bond
Periodic Trend:
Electronegativity decreases down a group
Electronegativity increases across a row (left to right)
Fluorine is the most electronegative atom
Cesium atom has a low electronegativity