Chapter 12 The Periodic Tableblanchettechemlab.weebly.com/.../37713227/periodic_table.pdfHistory...
Transcript of Chapter 12 The Periodic Tableblanchettechemlab.weebly.com/.../37713227/periodic_table.pdfHistory...
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The Periodic Table
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History
1829 German J. W. Dobereiner
Grouped elements into triads
One of these triads included chlorine, bromine, and
iodine; another consisted of calcium, strontium, and
barium. In each of these triads, the atomic weight of
the intermediate element is approximately the
average of the atomic weights of the other two
elements. The density of that element is
approximately the average of the densities of the
other two elements.
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History
The problem with this arrangement was that
Dobereiner’s model became outdated as new
elements were identified.
A good model is able to incorporate newly
understood information.
Dobereiner’s Triad Model was not useful,
since several newly discovered elements did
not “fit” into it.
Not all elements had triads
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History
Russian scientist Dmitri Mendeleev
taught chemistry in terms of
properties
Mid 1800 – atomic masses of
elements were known
Wrote down the elements in order of
increasing atomic mass
Found a pattern of repeating
properties
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Mendeleev’s TableGrouped 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 TableElements 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
There are 7 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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1A
2A 3A 4A 5A 6A7A
8A
0
The elements in the A groups
are called the representative
elements
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Metals
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Metals Luster – shiny.
Ductile – drawn into wires.
Malleable – hammered into sheets.
Conductors of heat and electricity.
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Transition metals The Group B
elements
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Non-metals Dull
Brittle
Nonconductors
- insulators
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Metalloids or Semimetals Properties of both
Semiconductors
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These are called the inner
transition elements and they
belong here
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Group 1A are the alkali metals
Group 2A are the alkaline earth metals
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Group 7A is called the Halogens
Group 8A are the noble gases
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Group Characteristics Alkali metals
• Group 1: very reactive metals which do not occur freely in nature.
1 electron in outer shell
Alkaline Earth Metals
• Group 2: next reactive metals, found in earths crust but not in elemental
form.
2 electrons in outer shell
Transition Elements
• Group 3-12: metals with varying reactivities. Greater density than Group 1
or 2 elements.
1-2 electrons in outer shell
Lanthanides and Actinides
• These elements are also transition elements but have been taken out to
prevent the periodic table being so wide.
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Boron Group
• Group 13: reactive, contains metal and metalloid.
3 electrons in outer shell
Carbon Group
• Group 14: contains metalloids, metals and non metals.
4 electrons in outer shell
Nitrogen Group
• Group 15: contains metalloids, metals and non metals.
5 electrons in outer shell
Oxygen Group
• Group 16: contains contains metalloids, metals and non metals. Reactive
6 electrons in outer shell
Halogens
• Group 17: non-metals, very reactive.
7 electrons in outer shell
Nobel gas
• Group 18: non-metals, non reactive.
8 electrons in outer shell
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Periodic Table Part 2
Periodic trends
Identifying the patterns
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What we will look for
Periodic trends-
• How properties vary as you go across
a period
Group trends
• How properties vary as you go down
a group
Why?
• Explain why they vary
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What we will investigateAtomic size
• how big the atoms are
Ionization energy
• How much energy to remove an electron
Electronegativity
• The attraction for the electron in a compound
Ionic size
• How big ions are
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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.
Increased atomic radius decreases IE
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Group trendsAs you go down a group first IE
decreases because of
• Bigger atoms…so outer electron
less attracted even though there
are more +charges in the nucleus
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Periodic trends
All the atoms in the same period
• Same approximate size
• Increasing nuclear charge
So IE generally increases from left to
right.
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Discussion Question
What is Atomic Radius?
• The radius of an atom
• The distance an atom travels
• The width of an atom
• The distance around an atom
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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 molecule}
Radius
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Trends in Atomic Size
Influenced by two factors
• Energy Level (Electron Shell)– Higher energy level (shell) is further away
• Charge on nucleus– More charge pulls electrons in closer
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Group trendsAs we go down a
group each atom
has another
energy level
(electron shell)
So the atoms get
bigger
H
Li
Na
K
Rb
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Periodic TrendsAs you go across a period the radius
gets smaller.
Same energy level (electron shell)
More nuclear charge
Pulls outermost electrons closer
Na Mg Al Si P S Cl Ar
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What is Ionic Radius?
• the radius of an atom
• the radius of the most common ion of
an atom
• the size of an atom
• the width of an ion
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Ionic Size
Cations are positive ions
• Cations form by losing electrons
• Metals form cations
Cations are smaller than the atom
they come from
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Ionic size
Anions are negative ions
• Anions form by gaining electrons
• Nonmetals form anions
Anions are bigger than the atom they
come from
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Group trends
Adding energy level
(electron shell)
Ions get bigger as
you go down
Li1+
Na1+
K1+
Rb1+
Cs1+
H1+
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Periodic Trends
Across the period nuclear charge
increases so they get smaller.
Energy level (electron shell) changes
between anions and cations
Li1+
Be2+
B3+
C4+
N3-
O2- F1-
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Electronegativity
The tendency for an atom to attract
electrons to itself when it is
chemically combined with another
element.
• How “greedy”
Big electronegativity means it pulls
the electron toward it.
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Group Trend
The further down a group
• Bigger atoms (outer electrons further
from nucleus)
• More electrons per atom
Less attraction for electrons
Lower electronegativity.
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Periodic Trend
Metals - left end
• Low nuclear charge
• Low attraction for extra electrons
Low electronegativity
Right end - nonmetals
• High nuclear charge
• Large attraction for extra electrons
High electronegativity
Not noble gases- no compounds
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Ionization energy, electronegativity
INCREASE
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Atomic size increases,
Ionic size increases
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Nuclear Charge
Energ
y L
evels
& S
hie
ldin
g
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How to answer why questions
Trend
• Periodic and Group
Reason
• Nuclear charge
• Energy level and distance from
nucleus
Result
• What happens to which electron