1. TRENDS
IN THE
PERIODIC TABLE

3. ATOMIC SIZE
- Estimated by measuring the distance between 2 neighboring atoms

4. FACTORS AFFECTING ATOMIC SIZE
PRINCIPAL QUANTUM NUMBER (n) OF THE VALENCE ELECTRON – the farther the valence electron is from the nucleus, the greater is the atomic size.
EFFECTIVE NUCLEAR CHARGE (Zeff ) – charge “felt” by the valence electron due to the attractive force of the protons (Z) and the shielding effect (S) by the core electrons
Zeff = Z - S

5. Group trends As we go down a group...H
As we go down a group...
each atom has another energy level,
so the atoms get bigger. Li Na K Rb

6. Periodic Trends As you go across a period, the radius gets smaller.
Electrons are in same energy level.
More nuclear charge.
Outermost electrons are closer. Na Mg Al Si P S Cl Ar

7. Periodic Trends As you go across a period, the radius gets smaller.
Electrons are in same energy level.
More nuclear charge.
Outermost electrons are closer. Na Mg Al Si P S Cl Ar

8. Rb K
Overall Na Li
Atomic Radius (nm) Kr Ar Ne H 10
Atomic Number

9. +1-0=+1
+2-0=+2
+4 -2=+2
+3 -2=+1

11. EFFECT OF CHARGE ON SIZE OF IONS

12. IONIZATION ENERGY (IE) - Amount of energy needed to remove an electron from an atom.

13. Ionization Energy
The energy required to remove (1 mole of) the first electron is called the first ionization energy.
The second ionization energy is the energy required to remove (1 mole of) the second electron(s).
Always greater than first IE.
The third IE is the energy required to remove a third electron.
Greater than 1st or 2nd IE.

14. Symbol First Second Third
HHeLiBeBCNO F Ne

15. Symbol First Second Third
HHeLiBeBCNO F Ne

17. What determines IE The greater the nuclear charge, the greater IE.
Greater distance from nucleus decreases IE
Filled and half-filled orbitals have lower energy, so achieving them is easier, lower IE.
Shielding effect

18. Shielding
The electron in the outermost energy level experiences more inter-electron repulsion (shielding).
Second electron has same shielding, if it is in the same period

19. Group trends As you go down a group, first IE decreases because...
The electron is further away.
More shielding.

20. Periodic trends
All the atoms in the same period have the same energy level.
Same shielding.
But, increasing nuclear charge
So IE generally increases from left to right.
Exceptions at full and 1/2 full orbitals.

21. He has a greater IE than H. same shielding greater nuclear charge
First Ionization energy
Atomic number

22. Outer electron further away outweighs greater nuclear charge
Li has lower IE than H
Outer electron further away
outweighs greater nuclear charge H
First Ionization energy Li
Atomic number

23. greater nuclear chargeHe
Be has higher IE than Li
same shielding
greater nuclear charge
First Ionization energy H Be Li
Atomic number

24. greater nuclear chargeHe
B has lower IE than Be
same shielding
greater nuclear charge
p orbital is slightly more diffuse and its electron easier to remove
First Ionization energy H Be B Li
Atomic number

25. First Ionization energyHe
First Ionization energy H C Be B Li
Atomic number

26. First Ionization energyHe N
First Ionization energy H C Be B Li
Atomic number

27. First Ionization energyHe
Breaks the pattern, because the outer electron is paired in a p orbital and experiences inter-electron repulsion. N
First Ionization energy H C O Be B Li
Atomic number

28. First Ionization energyHe F N
First Ionization energy H C O Be B Li
Atomic number

29. Ne has a lower IE than He Both are full, Ne has more shielding
Greater distance F N
First Ionization energy H C O Be B Li
Atomic number

30. Na has a lower IE than Li Both are s1 Na has more shieldingHe Ne
Na has a lower IE than Li
Both are s1
Na has more shielding
Greater distance F N
First Ionization energy H C O Be B Li Na
Atomic number

31. First Ionization energy
Atomic number

32. Driving Force
Full Energy Levels require lots of energy to remove their electrons.
Noble Gases have full orbitals.
Atoms behave in ways to achieve noble gas configuration.

33. ELECTRON AFFINITY (EA) – amount of energy needed to add electrons to an atom

35. Trends in Electron Affinity
The energy change associated with adding an electron (mole of electrons) to a (mole of) gaseous atom(s).
Easiest to add to group 7A.
Gets them to full energy level.
Increase from left to right: atoms become smaller, with greater nuclear charge.
Decrease as we go down a group.

36. ELECTRONEGATIVITY – tendency of an atom to attract electrons; guided by the octet rule.