HONORS CHEMISTRYSection 5.3 – Electron Configuration and Periodic Properties

Atomic Radius

Determined by the distance from the nucleus to the edge of the outer orbital.

Edge of outer orbital not well defined Use identical bonded atoms – ½ of

the distance between the nuclei

Trends in Atomic Radii

Trends in Atomic Radii

Decrease across the period Due to increasing positive charge of the

nucleus Increase as you go down the group

Exception Ga to Al – Ga smaller due to increased nuclear charge (first addition of d electrons)

Problems

Which of these elements; Li, Rb, K or Na has the smallest radius? Largest?

Which of these elements; O, Se, S and Po has the smallest atomic radius? Largest?

Ionization Energy

Atom + energy → Atom+ + e-

First electron removed – First Ionization Energy (IE1)

Second electron removed – Second Ionization Energy (IE2) etc.

Group 1 – lowest ionization energy Group 18 - highest ionization energy Ionization energies increase across the period due

to increased nuclear charge. Ionization energies decrease down a group due to

further distance from nucleus and electron shielding

Ionization Energies

Ionization Energies

Successive Ionization Energies

Why?

Each successive electron feels a stronger nuclear attraction

This information lead to the understanding of the stability of the noble gas configuration+

Practice

Choose the element with the higher IE1: Ca or Ba Ca or Br Ca or K Ca or Mg

Electron Affinity

Atom + e- → Atom- + (- energy)

IMPORTANT!!!!! Negative energy means energy lost by system Positive energy means energy gained by

system Sign indicates direction not numerical value!!!!

Electron Affinity Values

Electron Affinity Trends

Generally become larger (look at as absolute value) as you move across the period.

Exception – Group 15 due to half filled p orbitals

Generally become smaller as you move down a group due to: Greater Nuclear Attraction Greater Atomic Radius

Second Electron Affinities Very difficult to add an electron to an

anion (negative ion) Second Electron Affinities are all

positive

Ionic Radii

Cation (positive ion) Smaller atomic radius than atom Due to:

electrons being removed increased effective nuclear charge

Anion Larger atomic radius than atom Due to

electrons being added decreased effective nuclear charge Greater repulsion of electrons

Ionic Radii

Valence Electrons

Available to be lost, gained or shared when compounds are formed.

In outer main energy levels For Main Group Elements – s and p

orbitals

Bonded Atoms

Very rarely are electrons shared equally

Usually attracted more to one atom This will effect the chemical

properties of the compound!!! Measure of attraction – called

electronegativity Based on a 4.0 scale – F = 4.0. Developed by Linus Pauling

Electronegativity

Electronegativity Trends

Increase across a period. Tend to decrease or stay the same

down a group. If a noble gas does not form

compounds – it does not have an electronegativity

If a noble gas does form compounds – it will have a high electronegativity

Summary of Trends

Summary of Trends

D-Block

These elements tend to vary less and with less regularity than Main Group Elements.

Still electrons in d orbitals are often responsible for characteristics of elements in the d-block

Atomic radius tends to decrease across the block

Ionization energies generally increase across both the d and f-blocks

D and F Blocks

Tend to lose electrons from outer shell!!!!

That means the valence electrons come from the ns shell not the (n-1)d shell

Generally these elements from 2+ ions.

Electronegativities D-block - between 1.1 and 2.54 (Only

groups 1 and 2 are lower) F-block – between 1.1 and 1.5