Periodic Table: Patterns

John Newlands

• 1864

• arranged elements in octaves

• worked for some elements, but not all

Dimitri Mendeleev

• Julius Lothar Meyer

• both independently created a form of the periodic table

Dimitri Mendeleev

• Mendeleev usually credited because he showed how useful the table could be to predict properties

Using the periodic table

• The periodic table has so much information contained in its organization that it will become your most valuable resource!!!

Periodic Law

• Elements are arranged by atomic number, there is a periodic repetition in their physical and chemical properties.

Periodic Table

• Horizontal rows: – Called “periods”– 7 periods

• Show the number of the shell

Periodic Table

• Vertical columns: – Called “groups”– Show the valance electrons– Elements in a group have similar physical and

chemical properties

Periodic Trends

• We have already learned one periodic trend-

• the way electrons are organized in atoms

Some exceptions

• Draw the electron configuration of:

• Cr

• Cu

Single electron

• The energy of a single electron reflects how tightly bound that e- is to the nucleus (more negative energy)

Effective Charge

• We assume that the electrons are bound by a positive charge Z

• But there is an effective charge that e- “feels”

Effective Charge

• The effective charge (Zeff) represents the net effect of the attraction of the nuclear charge and the repulsion of the other e-’s

Effective Charge

• An electron closer to the nucleus would feel more of the positive charge

Shielding

• The “protection” by the inner electrons so that the outer electrons do not feel the full nuclear charge

Orbital Filling

• Shielding helps explain why the orbitals are filled in the order that they are

• The lower energy orbitals (ones closer to the nucleus) are always filled first

Orbital Shielding

• Orbitals to the inside of the other orbitals do a good job of shielding the outer electrons

Polyelectronic Model

• The energy required to remove an electron from an atom depends on two factors:

Two Factors

• The effective nuclear charge

• The average distance of the electron from the nucleus

Atomic radiusHalf the distance between the nuclei of two atoms. (Atoms are the same and bonded together)

Element Characteristics

Radius size decreases

Radius size

increases

Element Characteristics• Ionization energy

– Energy required to remove an electron from an atom • Easier to remove an electron from group 1 than

group 8• Easier to remove electrons that are farther away

from the nucleus (elements at the bottom of a group)

Ionization energy increases

Ionization energy

generally decreases

The general trend

• As we go across a period from left to right, the first ionization energy increases

Why??

• Electrons added in the same principle quantum level do not completely shield the increasing nuclear charge

Exceptions

• Decreases in ionization energy moving across (i.e. N to O) is because e- in 2s provide some shielding for the 2p

Electron Affinity

• The energy change associated with the addition of an electron to a gaseous atom

• X(g) + e- --> X-(g)

Sign for the energy

• Defined as energy change when electron is added

• if addition is exothermic then the sign is negative

Electron affinity

• What is the trend going down a group?

• Generally becomes more positive b/c e- added at increasing distance

Trends...

• Electron affinities generally become more negative from left to right across a period, there are several exceptions

Element Characteristics• Electronegativity

– It is a measure of the element's ability to attract the electrons which are in a bond

Electronegativity increases

Electronegativity decreases

Element Characteristics

Liquids

Solids More metallic

More nonmetallic Periodic

Table

Periodic Table Trends: Summary

Atomic radius decreases

Ionization energy and

Electronegativity

increase