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The Periodic Table
Unit 4 Page 4
Learning Target:
I can use the Periodic Table to identify and explain periodic trends, including atomic and ionic radii, electronegativity,
and ionization energy.
Criteria For Success:
I can explain effective nuclear charge and electron shielding.
I can explain how increasing effective nuclear charge across a period influences the atomic radius, electronegativity and
ionization energy.
I can explain how increased electron shielding down a group influences atomic radius, electronegativity and ionization
energy.
I can explain the trend of ionic radius and how the magnitude and the sign of the ionic charge influences ionic radius
Notes: Periodic Trends
A. Attraction between the nucleus and electrons
1. Coulomb’s law- the force of attraction between two _________________ charged particles is
______________ proportional to the magnitude of the charges and __________________ proportional to the
distance between those charges.
𝑘𝑞1 𝑞2
𝐹=
𝑟2
2. Effective nuclear charge (Zeff)- the net _______________ _________ experienced by the _______________
electrons in an atom.
a. Zeff = # protons - # core (non-valence) electrons
3. Electron shielding- the inner electrons ______________ the outer electrons from some of the nucleus’
___________________ force.
a. Going down a group adds another ___________________ energy level  adds more shielding.
B. Periodic trends are _________________ in the periodic table relating to the ________ and other ________________
of elements.
1. ATOMIC RADIUS is the ________________ between the nucleus and the ____________ edge of the electron
cloud.
2. IONIZATION ENERGY is the _____________ required to _________________ an electron from the atom.
a. Each additional electron requires ___________ energy to remove than the previous one, so the
________________ ionization energy will be greater than the __________ ionization energy, and so on.
c. EXCEPTIONS:
i. O vs. N- It’s easier to remove an electron from O because of the ________________ of the
paired electrons.
ii. Be vs. B- It’s easier to remove an electron from B because it’s a __________________ vs. an
____________________ and electrons in p orbitals are held less _________________.
3. ELECTRONEGATIVITY is the ability of an atom in a __________________ [meaning it’s participating in a
___________] to attract _________________ electrons to itself.
a. Think of the atoms as playing “tug of war” with their ____________ shell electrons.
4. IONIC RADIUS is the ________________ from the nucleus to the outer edge of the electron cloud in a
________________ ion.
a. Cation- the atom __________ electron(s), has a ________________ charge
b. Anion- the atom ______________ electron(s), has a ________________ charge
c. All atoms want a __________ valence shell, and will gain/lose electrons to get there.
The Periodic Table
Unit 4 Page 5
Draw arrows representing the trend of each of the following from decreasing to increasing: Electronegativity, atomic
radius, Ionizations energy, electron affinity, reactivity
1. Organize the following elements in order of decreasing to increasing for each property: Mg, S, Fr, F, Al, Zn
Electronegativity:
Atomic radius:
Ionization energy:
Electron affinity:
2. What happens to atomic radius as you go across a period? Why does this happen?
3. What happens to reactivity as you go down a group? Why does this happen? Organize the following in
increasing reactivity: Be, Ca, Ba, Mg, Ra, Sr.
4. Why does an electron increase in electronegativity as it moves across the periodic table?
5. How does electron shielding contribute to the trend of increasing reactivity as you go down a group for metals?
How does it contribute the trend of increasing reactivity as you go up a group for non-metals?