Periodic Relationships Among the Elements

Chapter 8

Dr. Ali Bumajdad

Chapter 8 Topics

•History of periodic table and periodic trends•Periodic Classification of element•Periodic trend of Physical properties (Atomic size, Ionization energy, Electron Affinity, and Electronegativity )

Ions electron configuration •Ions Sizes (isoelectronic ions)•Group elements and its oxides

Periodic Relationships

History of periodic table and periodic trends

•Lother Meyer (German, 1830-1895) and Ivanonich Mendeleev(Russian, 1834-1907) put the form of periodic table

• Mendeleev given most credit because he predict the existence and properties of unknown elements and he corrected several values for atomic masses

Mendeleev’s periodic Table

Corrected (they thought its atomic mass is 76)

Atomic massNot atomic number

When the Elements Were Discovered

The current periodic table is different than Mendeleev by:

1) Element order by atomic number not by atomic mass

2) Contain many more elements discovered after Mendeleev.

Classification of the Elements

•Atomic radius•Ionization energy•Electron affinity•Electronegativity

•Atomic radius•Ionization energy•Electron affinity•Electronegativity

-Electron is added ondifferent principle quantum numberEffective nuclear +ve charge not much change due to the shielding effect

Size increase

-Electron is added on same principle quantum numberEffective nuclear +ve charge increase due to no shielding effect Size decrease

(1) Atomic size (2) Ionization energy (3) Electron affinity

•Atomic radius is half thedistance between twoneighboring atoms in their solid state

•As you go down a groupthe atom size increase

(Zeff is approx. constant

or even decrease and e placed in a new energy level)

•Energy required to remove e from an isolated gaseous atom, ion or molecules (kJ/mol)

Na(g) + E Na+ + 1e

• Endothermic process (+ve E) because e is attracted to nucleus

• As size increase it is easier to remove e

•Effective nuclear charge (Zeff): the “positive charge” felt by an electron

•Energy released or absorbed when an electron is added to an isolated gaseous atom, ion (kJ/mol)

Cl-(g) + 1e + E Cl2-

Cl(g) + 1e Cl- + E

Release =exothermic

Absorbed = endothermic

(1) Atomic size (2) Ionization energy (3) Electron affinity

•As you go left to right across a period atom get

smaller (Zeff increase and e

placed in the same energy Level.

•For transition elements and inner transition elements, the last e,s are added ontoan inner shell (3d) not the valence shell. This cause

the Zeff to increase but not

much (there is a slight shielding effect) atomic size decrease but slowly.

• All element (except H) have more than 1 ionization energy

•1st Eionization < 2ed Eionization <…

(removing e from neutral atom is easier than from +ve ion)

Q) Why the 2nd Eionization for alkali metal is much more than the first one?

It has a noble gas electron configuration

• 1st Eaffinity usually exothermicbecause e placed into environment where it experience attraction to nucleus

• All element have more than 1 Electron Afffinity

• 2st Eaffinity endothermicbecause the second e added to a negatively charge ion

(1) Atomic size (2) Ionization energy (3) Electron affinity

Q) Why Zr have similar size as Hf?

A raw of Lanthanide elements

•For ions:

+ve ion < -ve ion

Q) Why the 3ed Eionization for earth alkali metal is the highest change?

It has a noble gas electron configuration

• Some Irregularities a cross a period

P Eaffinity < As affinity

O Eaffinity < S Eaffinity

Why?•Because e are crowding in shell 2 more than in shell 3more repulsion in shell2less energy evolveless Eaffinity

1

2 C has -ve Eaffinity while N has +ve Eaffinity

(1) Atomic size (2) Ionization energy (3) Electron affinity

• Some Irregularities a cross a period

•For Be and B, it is easier to remove e from 2p1 than from 2s2 because it is further apart from nucleus

Be Eionization > B Eionization

N Eionization > O Eionization

Why?

•For N and O, it is easier to remove e when there is e-e repulsion in an orbital

P Eionization > S Eionization

•Because the e in c inter a vacant 2p orbital while for N e must be placed in an already occupied orbital

Effective Nuclear Charge (Zeff)

increasing Zeff

incr

easi

ng Z

eff

Adding electron to the same principal quantum number

Adding electron on different principal quantum number

Be2+ < Mg2+ < Ca2+ < Sr2+

Sa Ex. : Eionization of P = 1060kJ/mol and for S = 1005 kJ/mol, explain this irregularity.

The valence electron configuration for S contains one filled orbital and hence there is e-e repulsion and hence easier to be removed

Sa Ex. : Among the following elements (Ne, Na, Mg) (a) Which atom has the largest 1st ionization(b) Which atom has the smallest 2nd ionization energy

(a) Ne(b) Mg (the first and second e’s are valence electrons)

The Alkali Metals

Li

Cs •More reactive (why)

• In aqueous system less reducing ability (why)Difficult to be solubilized in water due to its large size (less effective nuclear charge), less polar

Easier to loss electron (less effective nuclear charge)

Na [Ne]3s1 Na+ [Ne]

Ca [Ar]4s2 Ca2+ [Ar]

Al [Ne]3s23p1 Al3+ [Ne]

Atoms lose electrons so that cation has a noble-gas outer electron configuration.

H 1s1 H- 1s2 or [He]

F 1s22s22p5 F- 1s22s22p6 or [Ne]

O 1s22s22p4 O2- 1s22s22p6 or [Ne]

N 1s22s22p3 N3- 1s22s22p6 or [Ne]

Atoms gain electrons so that anion has a noble-gas outer electron configuration.

Ions electron configuration•When two nonmetals react they share electrons, so that the valence electron configuration completes for both atoms

•When nonmetal react with metal, the ions form so that the valence electron configuration of the nonmetal achieves the electron configuration of the next noble gas atom and the valence orbital of the metal are emptied

+1

+2

+3 -1-2-3

Cations and Anions Of Representative Elements

Electron Configurations of Cations of Transition Metals

When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals.

Fe: [Ar]4s23d6

Fe2+: [Ar]4s03d6 or [Ar]3d6

Fe3+: [Ar]4s03d5 or [Ar]3d5

Mn: [Ar]4s23d5

Mn2+: [Ar]4s03d5 or [Ar]3d5

Na+: [Ne] Al3+: [Ne] F-: 1s22s22p6 or [Ne]

O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne]

Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne

Dr. Ali Bumajdad

Ions Sizes (isoelectronic ions)

•Isoelectronic ions: ions contain the same number of electrons

No. proton 11 13 9 8 7 10

The ion with less protons is bigger

N3- > O2- > F- > Ne > Na+ > Al3+

Q) What neutral atom is isoelectronic with H- ?

H-: 1s2 same electron configuration as He

Sa. Ex. : Arrange the ions in order of decreasing size ?Se2-, Br-, Rb+, Sr2+

Se2- > Br- > Rb+ > Sr2+

Sa. Ex. : Choose the largest ion in each of the following Groups:a) Li+, Na+, K+, Rb+, Cs+

b) Ba2+, Cs+, I-, Te2-

Cation is always smaller than atom from which it is formed.Anion is always larger than atom from which it is formed.

Dr. Ali Bumajdad

The Radii (in pm) of Ions of Familiar Elements

Sa Ex. : Predict the trend in radius for : Be2+ ,Mg2+ ,Ca2+ ,Sr2+

Group 1A Elements (ns1, n 2)

Group 2A Elements (ns2, n 2)

Group 3A Elements (ns2np1, n 2)

Group 4A Elements (ns2np2, n 2)

Group 5A Elements (ns2np3, n 2)

Group 6A Elements (ns2np4, n 2)

Group 7A Elements (ns2np5, n 2)

Group 8A Elements (ns2np6, n 2)

8.6

Completely filled ns and np subshells. Highest ionization energy of all elements.No tendency to accept extra electrons.

Properties of Oxides Across a Period

basic acidic