1. Atomic Radius

2. What it is: Measure of the size of the atom
Distance from nucleus to the boundary of the surrounding cloud of electrons
Difficult to define because the electrons do not have defined orbits so their positions are estimated.

3. Periodic Trends (Going Down a Column)
Principle quantum number increases
Atomic radius increases
examples:
F = .71?
Cl = .99?
Br = 1.14?
I = 1.33?

4. Periodic Trends (Left to Right in a Period)
Effective nuclear charge increases
Atomic radius decreases
examples:
Na = 1.86?
Al = 1.43?
S = 1.04?
Ar = .97?

5. How to measure: No fixed radius
Measure the distance between the nuclei of two bonded atoms
The radii is determined by the bonds that they form

6. Example Problem:
The atomic radius of F=64, Br=114 and I=138 pm respectively. Based on this information estimate a reasonable atomic radius of Cl.
53 pm
89 pm
126 pm
162 pm
196 pm

7. Example Problem:
Using the periodic table and the known periodic trends, which element has the largest atomic radius? H He Fr Rn Rh

8. Ions and Ionic Radii: cation anion positive charge negative charge
smaller atomic radius than the uncharged atom
example:
Na = 191
Na + = 95
Li = 156
Li + = 60
anion
negative charge
larger atomic radius than the uncharged atom
example:
F = 62
F- = 133
Cl = 102
Cl- = 181

9. Periodic Trends of Ionic Radii:
Decreases across the period (from left to right)
Increases within groups (from top to bottom)

10. Metals & NonMetals Metals: NonMetals: decreases after it ionizes
metals lose electrons
NonMetals:
increase after it ionizes
nonmetals gain electrons

11. Ionization Energy

12. Ion
Atom or molecule in which the total number of electrons is not equal to the total number of protons, giving the atom a net positive or negative electrical charge
Can be created by both chemical and physical means
All ions are charged: attracted to opposite electric charges, repelled by like charges, when moving, travel in trajectories that are deflected by a magnetic field

13. Trends of Ionization Energy
The energy that is needed to remove an electron from an atom is a process called ionization energy.
Ionization energy generally decreases from top to bottom within a group, while it increases from left to right across a period.
Ionization energy can be helpful when predicting what ions may be formed.
It is much easier to remove one electron, but when you go to remove a second electron the difficulty involved increases

14. First Ionization Energy
The amount of energy required to remove one electron.
For two electrons it is called the second ionization energy.
For the “S” block these two ionizations are useful.

15. Group Trends in Ionization Energy
Within each group (column): First ionization energy decreases from top to bottom.
For example, see the noble gases listed at each peak; they are in the same group.
Within each period (row): FIE increases when moving towards the right-hand side.
There are some exceptions, evident if we look at the elements from K-Kr.
This can be explained as the nuclear charge increases across the period.
At the same time, the shielding effect does not change.
Shielding effect: average amount of electron density between one e- and the nucleus
This inhibits the e-’s attraction to the nucleus.
Nuclear charge: Relative charge minus shielding effect constant
The higher the nuclear charge is the less shielding there is.

16. Visual Chart
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17. Shielding
In a multi-electron system, electrons are simultaneously attracted to the nucleus and repelled by each other
Outer electrons are shielded from nucleus by the core electrons
screening or shielding effect
outer electrons do not effectively screen for each other
The shielding causes the outer electrons to not experience the full strength of the nuclear charge

18. Effective Nuclear Charge
The effective nuclear charge is net positive charge that is attracting a particular electron
Z is the nuclear charge, S is the number of electrons in lower energy levels
electrons in same energy level contribute to screening, but very little so are not part of the calculation
trend is s > p > d > f
Zeffective = Z − S

19. Electronegativity The ability to gain electrons.
Generally, decreases from top to bottom
Generally, increases from left to right.
*Noble Gases