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How do you measure the size of an atom? The electron cloud doesn’t have a definite edge. Can get around this by measuring covalent atomic radius.
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Chemical Periodicity
Trends in the periodic tableTrends in the periodic table
Atomic Size
How do you measure the size of an atom?
The electron cloud doesn’t have a definite edge.
Can get around this by measuring covalent atomic radius.
Atomic Size
Atomic Radius = half the distance between two nuclei of a diatomic molecule.
}Radius
Atomic size is influenced by two factors:
Energy Level –more occupied levels = bigger atom
Charge on nucleus– More charge pulls electrons in closer
Group trends
As we go down a group electrons are added to higher energy levels so the atoms get bigger.
HLi
Na
K
Rb
Periodic Trends As you go across a period the radius
gets smaller. Same energy level, but protons pull
electrons closer to nucleus.
Na Mg Al Si P S Cl Ar
Trends in Atomic Radius
Questions:
Of the following elements, which has the largest atomic radius? Why?
a) Si, Mg, S
b) Al, Na, Cl
c) Li, Cs
Mg – same energy level, smallest nuclear charge
Na – same energy level, smallest nuclear charge
Cs – higher occupied energy levels
Ionic Size
Ionic SizeCations Positive ions - form by losing electrons. Metals form cations Cations of representative elements have noble gas
configuration. Smaller than the atom they come from because of
increased attraction by nucleus for fewer remaining electrons
+
Ionic sizeAnions Negative ions - form by gaining electrons. Nonmetals form anions. Anions of representative elements have noble gas
configuration. Larger than the atom they come from, because
nuclear attraction is less for an increased number of electrons.
-
Group trends
Ions get bigger as you go down (adding energy levels)
Li+1
Na+1
K+1
Rb+1
Cs+1
Periodic Trends
Across the period nuclear charge increases so both cations and anions get smaller from left to right.
Li+1
Be+2
B+3
C+4
N-3O-2 F-1
Questions:1. Of the following ions, which ones should have
the larger radius? Why?a) Na+ or Cs+
b) Br- or K+
2. The Mg2+ and Na+ ions have ten electrons surrounding the nucleus. Which ion would you expect to have the smaller radius? Why?
Cs+ It has more occupied energy levels
Br- Anions are larger than cations
Mg2+ Greater nuclear charge
Ionization Energy
Ionization Energy The amount of energy required to
completely remove an electron from a gaseous atom (how hard it is to pull an e- off an atom)
1st IE = removing 1 e-, 2nd IE=removing 2 e-
Na(g) Na+ + e-
Shielding
The electron on the outside energy level is shielded from the nucleus by the inner electrons
Group trends As you go down a group first IE
decreases because the electron is further away (more shielding)
Periodic trends
All the atoms in the same period have the same energy level (same shielding).
As you go from left to right, nuclear charge increases so IE generally increases.
Questions:1. Which element in the following sets has the
lowest ionization energy and why?a) B, C, F
b) K, Na, LiB – same energy level, smallest nuclear charge
K – electron farther away, more shielding
Electron Affinity
Electron Affinity The energy given off when an electron is
added to an atom how much an atom ‘wants’ an electron
F(g) + e - F -(g)
Electron AffinityGroup trends Generally decreases as we go down a group
because shielding increases
Periodic trends Increases from left to right as atoms become
smaller with greater nuclear charge
Questions:1. Of the following elements, which ones should
have the higher electron affinity? Why?a) Se or Te
b) Calcium or ChromiumSe – smaller atom
Chromium – greater nuclear charge
Electronegativity
Electronegativity
The tendency for an atom to attract electrons to itself when it is chemically combined (BONDED) with another element.
Big electronegativity means it pulls the electron towards itself.
Group Trends The further down a group the farther the
electron is away from the nucleus and the more electrons an atom has.
More willing to share = low electronegativity
So as you go down a group electronegativity decreases
Periodic Trends As we go from left to right across the table,
electronegativity increases, because nuclear charge is increasing and electrons are held in more strongly
Metals have low electronegativity Non-metals have high electronegativities
(they win the electron tug-of-war)
Questions:1. Which element would you expect to have the
highest electronegativity? Why?
2. Put the following elements in order of increasing electronegativity: Na, P, Cl
F smallest nonmetal
Na, P, Cl