periodic trends

2. Atomic radius, Ionization energy,
Electronegativity

3.  Elements are arranged side by side in order of
increasing atomic number – makes it possible
to see commonalities among them.
 Vertical – Groups/Columns of the periodic
table
 Horizontal – Periods/Rows of the periodic
table

4.  The atomic radius is one-half of the distance
between the nuclei of two atoms of the same
element when the atoms are joined.

5. Trends in atomic
radius

6.  Atomic Radius Increases Down a group
◦ Greater energy levels and orbitals surrounding
nucleus
◦ Outer electrons are further away from nucleus
◦ Shielding effect occurs as we go down the group
◦ Shielding effect: The reduction of attractive force
between a nucleus and its outer electrons due to
blocking effect of inner electrons

8.  Atomic Radius Decreases across a period
 Energy level stays the same
◦ No shielding changes happen
◦ Increase in # of protons, therefore positive charge
increases
◦ Pulls valence electrons closer

9.  Atomic number
increases
 Energy level (orbitals)
increases
 Shielding occurs
 Decrease nuclear
charge
 Radius increases
 Atomic number
increases
 Same Energy level , no
shielding
 Increase in nuclear
charge,
 Pulls electrons in tight
and decrease radius
Down a group Across a period

10.  An ion
◦ is an atom or group of atoms that
has a positive or negative charge
◦ Cations
 An ion with a positive charge
 Metals
◦ Anions
 An ion with a negative charge
 Nonmetals

11. ◦ The energy required to
remove an electron from an
atom is called ionization
energy.

12.  ionization energy is high,
◦ takes a lot of energy to remove the outermost
electron
 ionization energy is low,
◦ takes only a small amount of energy to remove the
outermost electron.
 Smaller atomic radius, what do you think the
IE is? Why?
◦ High IE, small radius, high nuclear charge therefore
holding valence electrons tight

13.  Ionization energies can help you predict what
ions an element will form.
◦ It is relatively easy to remove one electron from a
Group 1A metal atom, but it is difficult to remove
a second electron.
◦ This difference indicates that Group 1A metals
tend to form ions with a 1+ charge

14. Trends of Ionization
Energy
Energy generally increases
Energy generally
decreases

15.  The ability of an atom to attract and hold an
extra electron
 - Measured as the energy change that occurs
when an electron is added to an atom
 - When an atom gains an extra electron it
becomes negative
 - Electron affinity can be either a positive or
negative numerical value
 - Negative Value: atom releases energy when it
gains an electron
 - Positive Value: atom must absorb energy for
the electron to be added

16.  Becomes More negative as we move from left
to right across a period
 Explained by nuclear charge, atomic radius
and shielding effects
 As we move down the groups, electron
affinity becomes less negative

17. ◦ Electronegativity
 the ability of an atom of an element to attract
electrons when the atom is in a compound.

18. The least electronegative element in
the table is cesium, with an
electronegativity of 0.7.
 It has the least tendency to attract electrons.
 When it reacts, it tends to lose electrons and form
cations.

20. The most electronegative element is
fluorine, with a value of 4.0.
 fluorine has such a strong tendency to attract
electrons
 when it is bonded to any other element it either
attracts the shared electrons or forms an anion.

21. In general, electronegativity
values decrease from top to
bottom within a group. For
representative elements, the
values tend to increase from left
to right across a period.
Metals at the far left of the periodic table have low values.
By contrast, nonmetals at the far right (excluding noble
gases) have high values.
Values among transition metals are not as regular.