1. Classification of Elements and Periodic Trends
CP Chemistry Section 6.3

2. Types of Electrons Valence electrons Core electrons
electrons filling the highest principal energy level of an atom
the outer electrons around an atom
involved in chemical reactions and contribute to an element’s chemical properties
found in s and p orbitals (maximum of 8 valence electrons)
Core electrons
electrons in lower orbitals close to the nucleus
are NOT involved in chemical reactions

3. Valence Electrons and the Periodic Table
One of the most important relationships in Chemistry: Atoms in the same group have similar chemical properties because they have the same number of valence electrons.
Valence electrons and period
Energy level of atom’s valence electrons = period
Ex. Elements in period 3 have valence electrons in energy level 3
Valence electrons and group number
Group number = number of valence electrons (applies only to A groups)
Ex. Group 3A elements all have 3 valence electrons

4. Orbital Blocks in the Periodic Table
Periodic table is divided into blocks based on which outer orbitals are being filled – can use them to predict an element’s electron configuration (especially end configurations)
s block = Groups 1A and 2A (2 groups)
p block = Groups 3A through 8A (6 groups)
d block = Transition metals in middle of table, B groups (10 groups)
f block = Inner transition metals (Lanthanide and Actinide series), found in two rows at bottom of table (14 columns)

5. Periodic Trends
Patterns in which properties of elements are related to their position in the periodic table
Trends exist for:
Atomic radius
Ionic radius
Ionization energy
electronegativity

6. Atomic Radius
Atomic size is defined by how closely an atom lies to a neighboring atom – because the electron cloud is based on the probability and is not clearly defined
Periodic trends
Move across period from left to right – atomic radii decrease, reason: each element across adds one more proton and electron – but the electron is added to the outer energy level so the nucleus holds onto electrons more tightly (no more electrons come between the nucleus and valence electrons)
Move down group from top to bottom – atomic radii increase, reason: each element in the family has more core electrons and valence electrons in higher energy levels, so the core electrons shield the valence electrons from the nucleus’ attraction to the valence electrons, therefore the radii increase

7. Trends in Atomic Radius

8. Ionic Radius Elements that lose valence electrons
Make positive ions
Hold their remaining electrons even more tightly
Have smaller radii than original atoms
Elements that gain valence electrons
Make negative ions
Newly added valence electrons spread out because of repulsion to other electrons
Have larger radii than original atoms

9. Trends in Ionic Radius

10. Octet Rule
Octet rule: atoms tend to gain, lose, or share electrons in order to acquire a full set of 8 valence electrons
Goal for non- Noble gas elements is to have a noble gas configuration (especially the representative elements)
Noble gases have a full octet – so very stable, do not react with other elements
Isoelectronic = having same electron configuration as another element (usually a noble gas)

11. Formation of Ions Based on number of valence electrons
Metals – have few valence electrons and lose them to form positive ions (cations)
Transition metals - can form more than one type of ion because they give away electrons from their s and d orbitals
Ex. Iron forms Fe2+ or Fe3+ ions
Configurations are not always isoelectronic but have special stability
Nonmetals – have many valence electrons and gain electrons to form negative ions (anions)

12. Formation of Ions – Groups 1A – 8A
Group 1A – loses one valence electron, +1 ions
Group 2A – lose two valence electrons, +2 ions
Group 3A – loses three valence electrons, +3 ions
Group 4A – half full valence: does not gain or lose, but shares electrons
Group 5A – gains three valence electrons, -3 ions
Group 6A – gains two valence electrons, -2 ions
Group 7A – gains one valence electron, -1 ions
Group 8A – full valence – does not gain, lose, or share electrons

13. Ionization Energy
Energy needed to remove an electron for a gaseous atom
First ionization energy = energy to remove the first electron from an atom
Indicates how strongly the atom holds onto its electrons
High ionization energy – atom has a strong hold on electrons, likely to form negative ions
Low ionization energy – atom has a weak hold on electrons, likely to form positive ions
Possible to remove more electrons after the first one – how much ionization energy required depends on if the electron is a valence electron or core electron
Atoms do not like to lose core electrons – will have very high ionization energies

14. Ionization Energy

15. Periodic Trends for Ionization Energy
Across Periods
Moving left to right, ionization energy increases
Reason: atoms are smaller and have more valence electrons so they hold onto them very tightly, becomes harder to remove electrons the closer the atom is to having a full valence
Down Groups
Moving from top to bottom, ionization energy decreases
Reason: atoms become larger and there is more shielding of the valence electrons by the core electrons (valence electrons are in higher energy levels and much farther from the attraction of positive charge in the nucleus)

16. Trend for Ionization Energy

17. Electronegativity
Indicates the relative ability of the atom to attract electrons in a chemical bond
Expressed in numerical values of 4 or less
Noble gases do not have electronegativities because they do not naturally form compounds
Show similar trends in periodic table as ionization energies

18. Periodic Trends for Electronegativity
Across Periods
Moving left to right, electronegativity increases
Reason: atoms with fuller valences have stronger pull on electrons within a bond
Down Groups
Moving from top to bottom, electronegativity decreases
Reason: atoms become much larger, have higher energy level valences, and due to shielding do not hold onto electrons in the bond as tightly

19. Trend for Electronegativity