Periodic Trends

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Periodic Trends

• Mrs.Kay

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• Groups vertical columns (18)
• Have similar properties because have same number
  of electrons in outer shell
• Periods horizontal columns (7)

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• Valence electrons electrons found in the outer
  most shell or valence shell
• Each energy level/shell holds on a certain number
  of electrons

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complicated element
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S- block
s1
s2

• Alkali metals all end in s1
• Alkaline earth metals all end in s2
• really have to include He but it fits better
  later.
• He has the properties of the noble gases.

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Transition Metals -d block
s1 d5
s1 d10
d1
d2
d3
d5
d6
d7
d8
d10
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The P-block
p1
p2
p6
p3
p4
p5
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F - block

• inner transition elements

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Metals To the left of the staircase line

• Chemical Properties
• Easily lose electrons
• Corrode easily (ex rusting or tarnishing)
• Low electronegativity

• Physical Properties
• Luster (shiny)
• Good conductors
• High density
• High melting point
• Malleable

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Non metals to the right of the staircase line

• Physical Properties
• Dull
• Poor conductor
• Brittle
• Not malleable
• Low density and melting point
• Chemical Property
• Tend to gain electrons
• High electronegativity

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Metalloids along the staircase line

• Solids
• Shiny or dull
• Malleable
• Conduct heat and electricity better than non
  metal but not as well as metals

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Atomic Size
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Atomic Size

• The electron cloud doesnt have a definite edge.
• They get around this by measuring more than 1
  atom at a time.
• Summary it is the volume that an atom takes up

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Atomic Size

Radius

• Atomic Radius half the distance between two
  nuclei of a diatomic molecule.

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Group trends
H

• As we go down a group (each atom has another
  energy level) the atoms get bigger, because more
  protons and neutrons in the nucleus

Li
Na
K
Rb
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Periodic Trends

• atomic radius decreases as you go from left to
  right across a period.
• Why? Stronger attractive forces in atoms (as you
  go from left to right) between the opposite
  charges in the nucleus and electron cloud cause
  the atom to be 'sucked' together a little tighter.

Na
Mg
Al
Si
P
S
Cl
Ar
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Reactivity
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Reactivity

• Reactivity refers to how likely or vigorously an
  atom is to react with other substances. This is
  usually determined by how easily electrons can be
  removed (ionization energy) and how badly they
  want to take other atom's electrons

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For Metals

• Period - reactivity decreases as you go from
  left to right across a period. Group -
  reactivity increases as you go down a group
• Why? The farther to the left and down the
  periodic chart you go, the easier it is for
  electrons to be given or taken away, resulting in
  higher reactivity

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For Non-metals

• Period - reactivity increases as you go from the
  left to the right across a period. Group -
  reactivity decreases as you go down the group.
• Why? The farther right and up you go on the
  periodic table, the higher the electronegativity,
  resulting in a more vigorous exchange of electron.

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Ionization Energy
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Ionization Energy

• The amount of energy required to completely
  remove an electron from a gaseous atom.
• An atom's 'desire' to grab another atom's
  electrons.
• Removing one electron makes a 1 ion.
• The energy required is called the first
  ionization energy.
• X(g) energy ?X e-

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Ionization Energy

• The second and third ionization energies can be
  represented as follows
•  X (g) energy? X2 (g) e-
• X2 (g) energy? X3 (g) e-
• More energy required to remove 2nd electron, and
  still more energy required to remove 3rd electron

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Group trends

• Ionization energy decreases down the group.
• Going from Mg to Be, IE decreases because
• Be outer electron is in the 3s sub-shell rather
  than the 2s. This is higher in energy
• The 3s electron is further from the nucleus and
  shielded by the inner electrons
• So the 3s electron is more easily removed
• A similar decrease occurs in every group in the
  periodic table.

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Period trends

• IE generally increases from left to right.
• Why?
•  From Na to Ar (11 protons to 18 protons), the
  nuclear charge in each element increases.
•  The electrons are attracted more strongly to
  the nucleus so it takes more energy to remove
  one from the atom.

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Why is there a fall from Mg to Al?

•  Al has configuration 1s2 2s2 2p6 3s2 3p1, its
  outer electron is in a p sublevel
•  Mg has electronic configuration 1s2 2s2 2p6
  3s2.
•  The p level is higher in energy and with Mg the
  s sub level is full this gives it a slight
  stability advantage

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Why is there a fall from P to S?

•  This can be explained in terms of electron
  pairing.
•  As the p sublevel fills up, electrons fill up
  the vacant sub levels and are unpaired.
• This configuration is more energetically stable
  than S as all the electrons are unpaired. It
  requires more energy to pair up the electrons in
  S so it has a lower Ionisation energy.
• There is some repulsion between the paired
  electrons which lessens their attraction to the
  nucleus.
•  It becomes easier to remove!

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Driving Force

• Full Energy Levels are very low energy.
• Noble Gases have full energy levels.
• Atoms behave in ways to achieve noble gas
  configuration.

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2nd Ionization Energy

• For elements that reach a filled or half filled
  sublevel by removing 2 electrons 2nd IE is lower
  than expected.
• Makes it easier to achieve a full outer shell
• True for s2
• Alkaline earth metals form 2 ions.

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3rd IE

• Using the same logic s2p1 atoms have an low 3rd
  IE.
• Atoms in the aluminum family form 3 ions.
• 2nd IE and 3rd IE are always higher than 1st IE!!!

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Electron Affinity
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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.

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Ionization energy, electronegativity Electron
affinity INCREASE
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Atomic size increases, shielding constant
Ionic size increases
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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.

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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.

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Configuration of Ions

• Ions always have noble gas configuration.
• Na is 1s22s22p63s1
• Forms a 1 ion 1s22s22p6
• 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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Periodic Trends

• Across the period nuclear charge increases so
  they get smaller.
• Energy level changes between anions and cations.

N-3
O-2
F-1
B3
Li1
C4
Be2
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Size of Isoelectronic ions

• Iso - same
• Iso electronic ions have the same of electrons
• Al3 Mg2 Na1 Ne F-1 O-2 and N-3
• all have 10 electrons
• all have the configuration 1s22s22p6

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Size of Isoelectronic ions

• Positive ions have more protons so they are
  smaller.

N-3
O-2
F-1
Ne
Na1
Al3
Mg2
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Electronegativity

• (optional coverage)

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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.
• So as you go from fluorine to chlorine to bromine
  and so on down the periodic table, the electrons
  are further away from the nucleus and better
  shielded from the nuclear charge and thus not as
  attracted to the nucleus. For that reason the
  electronegativity decreases as you go down the
  periodic table.

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Period Trend

• Electronegativity increases from left to right
  across a period
• When the nuclear charge increases, so will the
  attraction that the atom has for electrons in its
  outermost energy level and that means the
  electronegativity will increase

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Period trend

• Electronegativity increases as you go from left
  to right across a period.
• Why? Elements on the left of the period table
  have 1 -2 valence electrons and would rather give
  those few valence electrons away (to achieve the
  octet in a lower energy level) than grab another
  atom's electrons. As a result, they have low
  electronegativity. Elements on the right side of
  the period table only need a few electrons to
  complete the octet, so they have strong desire to
  grab another atom's electrons.

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Group Trend

• electronegativity decreases as you go down a
  group.
• Why? Elements near the top of the period table
  have few electrons to begin with every electron
  is a big deal. They have a stronger desire to
  acquire more electrons. Elements near the bottom
  of the chart have so many electrons that loosing
  or acquiring an electron is not as big a deal.
• This is due to the shielding affect where
  electrons in lower energy levels shield the
  positive charge of the nucleus from outer
  electrons resulting in those outer electrons not
  being as tightly bound to the atom.

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Shielding

• Shielded slightly from the pull of the nucleus by
  the electrons that are in the closer orbitals.
• Look at this analogy to help understand