Transcript Chapter 14 - Chemical Periodicity

Periodic Table
The most useful tool in the Lab
Early Organization
• J.W. Dobereiner (1829) organized elements in
triads
▫ Triad – three elements with similar properties (ex:
Cl, Br, I)
• J.R. Newlands (1864) organized elements in
octaves
▫ Octave – repeating group of 8 elements
Development of the PeriodiceTable
• Dmitri Mendeleev taught chemistry in
terms of properties.
• Mid 1800’s - molar masses of elements
were known.
• Wrote down the elements in order of
increasing mass.
• Found a pattern of repeating properties.
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.
• Also found gaps.
• Must be undiscovered elements.
• Predicted their properties before they were
found.
The Modern Periodic Table
• Henry Moseley – British physicist
• Arranged elements according to
increasing atomic number
• The arrangement today
• Symbol, atomic number & mass
The New Way
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•
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Elements are still grouped by properties.
Similar properties are in the same column.
Order is by increasing atomic number.
Added a column of elements (noble gases)
Weren’t found because they are unreactive.
Organization
• Horizontal rows = periods
▫ There are 7 periods
▫ Each period represents an energy level
▫ Every element in the same period has
 the same # of energy levels and
 the same core electron configuration
Organization
• Vertical column = group or family
▫ Similar physical & chemical prop.
▫ Same # of valence electrons
▫ Same common oxidation state
▫ Identified by number & letter
• Horizontal rows are called periods
• There are 7 periods
• Group 1A are the alkali metals
• Group 2A are the alkaline earth metals
• Group 7A is called the Halogens
• Group 8A are the noble gases
The group B are called the
transition elements
 These
are called the inner
transition elements, and they
belong here
• The elements in the A groups are
1A
called the representative elements
8A
2A outer s or p filling
3A 4A 5A 6A 7A
Lanthanides – the 4f orbital fills for these
elements
Actinide series – the 5f orbitals are being
filled for these elements.
Types of elements
• Metals
• Non-metals
• Metalloids or semi-metals
Metals
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•
•
•
•
•
•
Good conductor of heat and electricity
Malleable
Ductile
High tensile strength
High luster
Solid at room temperature
React by losing electrons
Nonmetals
• Poor conductors of heat and
electricity
• React by gaining electrons
• Some gases (O, N, Cl); some are
brittle solids (S); one is a fuming dark
red liquid (Br)
Semi-Metals
• Heavy, stair-step line
• Metalloids border the line
▫ Properties intermediate between
metals and nonmetals
▫ Learn the general behavior and
trends of the elements, instead of
memorizing each element property
• B, Si, Ge, As, Sb, Te
Families
Group IA – alkali metals
most reactive metals
Silvery in appearance
Soft
Combine easily with non-metals
Melting point is higher than the
boiling point of water
Have 1 valence electron
Families
• Group 2 – Alkaline Earth Metal Family
▫ Harder, stronger, denser, higher melting
point, and less reactive than alkali
▫ Usually not found as free elements, but
as compounds
▫ Have 2 valence electrons
Families
• Group 7 – Halogens
▫ Most reactive family
▫ Non-metals
▫ Have seven valence electrons
• Group 8 – Noble Gas
▫ Inert, unreactive
▫ Have full set of valence electrons
H
1
1s
1
Li
22s1
1s
3
Na
11
K
19
Rb
37
Cs
55
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p65s24d10 5p66s1
22s22p63s23p64s23d104p65s24d105p66s24f
1s
Fr
87 145d106p67s1
S- block
s1
s2
• Alkali metals all end in s1
• Alkaline earth metals all end in s2
• really should include He, but it fits
better later.
• He has the properties of the noble
gases.
1s2 He 2
Ne
2
2
6
1s 2s 2p
10
1s22s22p63s23p6 Ar
18
1s22s22p63s23p64s23d104p6 Kr
36
1s22s22p63s23p64s23d104p65s24d105p6 Xe
54
1s22s22p63s23p64s23d104p65s24d10
Rn
6
2
14
10
6
5p 6s 4f 5d 6p 86
The P-block
p1 p2
p3
p4
p5
p6
1
2
3
4
5
6
7
• Each row (or period) is the energy level for s
and p orbitals.
Areas of the periodic table
• Group A elements = s & p blocks
• representative elements
▫ Wide range of phys & chem prop.
Transition Metals -d block
1
d
2
d
3
d
s1
5
d
s1
5
6
7
8
10
10
d d d d d d
• d orbitals fill up after previous energy level,
so first d is 3d even though it’s in row 4.
1
2
3
4
5
6
7
3d
F - block
• inner transition elements
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
1
2
3
4
5
6
7
4f
• f orbitals start filling at 4f
5f
Atomic Size
}
Radius
• First problem: Where do you start measuring from?
• The electron cloud doesn’t have a definite edge.
•Atomic Radius = half the distance between two nuclei of
a diatomic molecule.
Trends in Atomic Size
Influenced by three factors:
1. Energy Level
Higher energy level is further away.
2. Charge on nucleus
More charge pulls electrons in
closer.
3. Shielding effect (blocking effect)
WHAT HAPPENS TO ATOMIC RADII?
• Does a negative ion (anion) get larger or smaller?
• Does a positive ion (cation) get larger or smaller?
Trends in Ionic Size
• Cations form by losing electrons.
• Cations are smaller than the atom
they come from.
• Metals form cations.
• Cations of representative elements
have noble gas configuration.
Ionic size
• Anions form by gaining electrons.
• Anions are bigger than the atom they
come from.
• Nonmetals form anions.
• Anions of representative elements
have noble gas configuration.
WHAT IS IONIZATION ENERGY?
• The energy required to remove an electron
• Which element has the highest ionization energy? Why?
What determines Ionization Energy?
• The greater the nuclear charge, the greater
IE.
• Greater distance from nucleus decreases IE
• All the atoms in the same period have the
same energy level.
• But, increasing nuclear charge
• So IE generally increases from left to right.
Ionization Energy
The energy required to remove the first
electron is called the first ionization
energy
The second ionization energy is the energy
required to remove the second electron.
Always greater than first IE.
The third IE is the energy required to
remove a third electron.
Greater than 1st or 2nd IE.
Driving Force
• Full Energy Levels require lots of
energy to remove their electrons.
• Noble Gases have full orbitals.
• Atoms behave in ways to achieve
noble gas configuration.
WHAT IS ELECTRONEGATIVITY?
• The ability of an atom to pull off an electron.
• Which element has the highest electronegativity?
Why?
Periodic Trend
• Metals are at the left of the table.
• They let their electrons go easily
• Low electronegativity
• At the right end are the nonmetals.
• They want more electrons.
• Try to take them away from others
• High electronegativity.
Trends in Electron Affinity
• The energy change associated with adding an
electron to a gaseous atom.
• Easiest to add to group 7A.
• Gets them to full energy level.
• Increase from left to right: atoms become
smaller, with greater nuclear charge.
• Decrease as we go down a group.