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Chapter 12The Periodic Table
The how and whyThe how and why
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History Russian scientist Dmitri Mendeleev
taught chemistry in terms of properties.
Mid 1800 - molar masses of elements were known.
Wrote down the elements in order of increasing mass.
Found a pattern of repeating properties.
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Mendeleev’s Table Grouped elements in columns by similar
properties in order of increasing atomic mass.
Found some inconsistencies - felt that the properties were more important than the mass, so switched order.
Found some gaps. Must be undiscovered elements. Predicted their properties before they
were found.
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The modern table Elements are still grouped by
properties. Similar properties are in the same
column. Order is in increasing atomic number. Added a column of elements Mendeleev
didn’t know about. The noble gases weren’t found because
they didn’t react with anything.
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Horizontal rows are called periods
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Vertical columns are called groups.
Elements are placed in columns by similar properties.
Also called families
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The group B are called the transition elements
These are called the inner transition elements and they belong here
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Group 1 are the alkali metals Group 2 are the alkaline earth metals
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Group 7 is called the Halogens Group 8 are the noble gases
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Why? The part of the atom another atom
sees is the electron cloud. More importantly the outside
orbitals. The orbitals fill up in a regular
pattern. The outside orbital electron
configuration repeats. The properties of atoms repeat.
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1s1
1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p65s24d10 5p66s1
1s22s22p63s23p64s23d104p65s24d105p66s2
4f145d106p67s1
H1
Li3
Na11
K19
Rb37
Cs55
Fr87
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He2
Ne10
Ar18
Kr36
Xe54
Rn86
1s2
1s22s22p6
1s22s22p63s23p6
1s22s22p63s23p64s23d104p6
1s22s22p63s23p64s23d104p65s24d105p6
1s22s22p63s23p64s23d104p65s24d10
5p66s24f145d106p6
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Atomic Size First problem where do you start
measuring. The electron cloud doesn’t have a
definite edge. They get around this by measuring
more than 1 atom at a time.
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Atomic Size
Atomic Radius = half the distance between two nuclei of a diatomic molecule.
}Radius
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Trends in Atomic Size Influenced by two factors. Energy Level Higher energy level is further
away. Charge on nucleus More charge pulls electrons in
closer.
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Group trends As we go down a
group Each atom has
another energy level,
So the atoms get bigger.
HLi
Na
K
Rb
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Periodic Trends As you go across a period the radius
gets smaller. Same energy level. More nuclear charge. Outermost electrons are closer.
Na Mg Al Si P S Cl Ar
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Overall
Atomic Number
Ato
mic
Rad
ius
(nm
)
H
Li
Ne
Ar
10
Na
K
Kr
Rb
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Ionization Energy The amount of energy required to
completely remove an electron from a gaseous atom.
Removing one electron makes a +1 ion.
The energy required is called the first ionization energy.
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What determines IE The greater the nuclear charge the
greater IE. Distance form nucleus increases IE Filled and half filled orbitals have
lower energy, so achieving them is easier, lower IE.
Shielding
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Shielding The electron on the
outside energy level has to look through all the other energy levels to see the nucleus
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Shielding The electron on the
outside energy level has to look through all the other energy levels to see the nucleus.
A second electron has the same shielding.
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Group trends As you go down a group first IE
decreases because The electron is further away. More shielding.
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Periodic trends All the atoms in the same period
have the same energy level. Same shielding. Increasing nuclear charge So IE generally increases from left to
right. Exceptions at full and 1/2 fill orbitals.
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Firs
t Ion
izat
ion
ener
gy
Atomic number
He
He has a greater IE than H.
same shielding greater nuclear
charge
H
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Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li has lower IE than H
more shielding further away outweighs greater
nuclear charge
Li
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Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Be has higher IE than Li
same shielding greater nuclear
charge
Li
Be
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Driving Force Full Energy Levels are very low
energy. Noble Gases have full orbitals. Atoms behave in ways to achieve
noble gas configuration.
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Ionic Size Cations form by losing electrons. Cations are smaller that the atom
they come from. Metals form cations. Cations of representative elements
have noble gas configuration.
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Ionic size Anions form by gaining electrons. Anions are bigger that the atom they
come from. Nonmetals form anions. Anions of representative elements
have noble gas configuration.
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Configuration of Ions Ions always have noble gas
configuration. Na is 1s12s22p63s1 Forms a +1 ion - 1s12s22p6 Same configuration as neon. Metals form ions with the
configuration of the noble gas before them - they lose electrons.
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Configuration of Ions Non-metals form ions by gaining
electrons to achieve noble gas configuration.
They end up with the configuration of the noble gas after them.
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Group trends Adding energy level Ions get bigger as
you go down.
Li+1
Na+1
K+1
Rb+1
Cs+1
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Electronegativity The tendency for an atom to attract
electrons to itself when it is chemically combined with another element.
How fair it shares. Big electronegativity means it pulls
the electron toward it. Atoms with large negative electron
affinity have larger electronegativity.
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Group Trend The further down a group the farther
the electron is away and the more electrons an atom has.
More willing to share. Low electronegativity.
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Periodic Trend Metals are at the left end. They let their electrons go easily Low electronegativity At the right end are the nonmetals. They want more electrons. Try to take them away. High electronegativity.
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Ionization energy, electronegativity
Electron affinity INCREASE
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Atomic size increases, shielding constant
Ionic size increases