Periodic Trends and
Bonding
Chapters 5 & 6
Ions and valence electrons
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How many valence electrons are in the
following elements?
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Na
Mg
H
He
Cl
Al
Show your Understanding
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Draw the Lewis dot structures for each of
those elements
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What ions will those elements form?
S- block
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Chemically reactive metals
Group #1= alkali metals
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Slippery appearance and can be cut with a knife!
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For real?
They all have one valence electron
Combine readily with the halogens to form salts
Group #2- Alkaline Earth Metals
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Harder, denser and stronger than group #1
metals
Have 2 valence electrons
d- block
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- Transition Metals
Lowest quantum # = 3
Maximum # of electrons = 10
There are exception to the electron
configuration rules
Some metals may form several different ions
They are all metals and good conductors of
heat and electricity and have high luster.
Properties vary greatly.
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Some metals are highly reactive
Other metals not so much- Au, Pt, Pd
p-block – Groups 13-18
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Group 13- 3 valence electrons
Group 14 – 4 valence electrons
Group 15- 5 valence electrons etc. etc.
Contains metals, non-metals and metalloids.
Important group- #17- Halogens
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Most are gases- most reactive with metals
Main Group elements
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Elements found in the s block and p block
Only elements that can be used in Lewis Dot
Structures
Key Concepts/ Terms
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Electrons have an attraction or pull towards
the nucleus of the atom.
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Shielding/ Screening: the attraction of outer
shell electrons is counterbalanced by the
repulsion of the inner-shell electrons. The
inner-shell electrons “screen” or “shield” the
outer-shell electrons from full attraction
Atomic Radii
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TREND:
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Increases from top to bottom
Decreases from left to right
Arrange the following in increasing atomic
radius:
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Rb, In, Sb, Sr, I
Ionization Energy
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The energy required to remove an electron
(IE2) = energy req. to remove a second electron
The more electrons removed from an atom the
greater the IE
TREND:
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Increases from left to right
Decreases from top to bottom
Exceptions: ( within the same energy level)
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Group 3A – lower energy than 1A and 2A bc entering p
orbital, slightly higher in energy than the s orbital for the
same level
Group 6A – electrons are paired up in the p-orbital and the
IE slightly dips, making it easier to remove
Ionization Energy
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Arrange the following elements in order of
decreasing IE.
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Na, Mg, Al, Si
Ionization Energy
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Arrange the following elements in order of
decreasing IE. Na, Mg, Al, Si
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Si < Al < Mg< Na
Electron Affinity
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Energy absorbed when an electron is added
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TREND:
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Greater Attraction for electrons from left to right
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( Left to right = least negative to most negative)
Decreases from top to bottom
Electron Affinity
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Which group would be the easiest to add an
electron to?
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Arrange the following elements in increasing
electron affinity:
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C, Si, Ge
Cl, S, P
Electron Affinity - answer
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Which group would be the easiest to add an
electron to?
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The halogens
Arrange the following elements in the most
negative to least negative electron affinity:
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C, Si, Ge – same
Cl, S, P – same
Ionic Radii- DO NOT DO
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Term: Isoelectronic : species that have the
same number of electrons. ( Na+, Mg2+)
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TREND:
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Top to Bottom = increases
Left to right = decreases
Ionic Radii
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Arrange the following ions in order of
increasing ionic radii:
Na+, Tl3+, Mg2+,
Ionic Radii - answer
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Arrange the following ions in order of
increasing ionic radii:
Mg2+, Na+, Tl3+,
Electronegativity
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Ability for an atom to attract electrons
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When it is chemically combined with another
atom.
Elements with high electronegativities
(nonmetals) often gain electrons to form
anions.
Elements with low electronegativities (metals)
often lose electrons to form cations.
Electronegativity
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TREND:
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Top to Bottom = decreases
Left to right = increases