1. Chemistry Chapter 6
The Periodic Law

3. Mendeleev’s Periodic Table
Dmitri Mendeleev

4. A Spiral Periodic Table
“Mayan” Periodic Table
Triangular Periodic Table

5. Giguere Periodic Table
Chinese Periodic Table
Stowe Periodic Table
Giguere Periodic Table

6. Orbital filling table

7. Periodic Law
Chemical properties are functions of an elements atomic number

8. Periodic Table with Group Names

9. Determination of Atomic Radius:
Half of the distance between nucli in
covalently bonded diatomic molecule
"covalent atomic radii"
Periodic Trends in Atomic Radius
Radius decreases across a period
Increased effective nuclear charge due
to decreased shielding
Radius increases down a group
Addition of principal quantum levels

10. Atomic Radius by period

11. Atomic Radius – Group He Be Mg Ca

12. Atomic Radius - Period
The positive charge in the nucleus increases with each new proton added – increasing the nuclear charge.
The negative charge on the outermost electron shell increases with each new electron added
The built-up nuclear charge attracts the built-up electron charge making the outer shell pull in closer to the nucleus as protons & electrons are added.

13. Electron Shells & Size

14. Check your knowledge

15. Ions
During chemical bonding atoms become stable by forming a noble gas electron configuration. Atoms will gain or lose electrons during bonding to achieve a noble gas electron configuration.
If an atom accepts or donates (gains or loses) electrons the atom will become an ion. An ion is a charged particle.
Positive ions are called cations.
Negative ions are called anions.

16. Ionization of Magnesium
Mg kJ  Mg+ + e-
1st ionization energy
Mg kJ  Mg e-
2nd ionization energy
Mg kJ  Mg e-
3rd ionization energy

17. Ionization Energy
Ionization Energy is the measure of energy in KJ/mole that are required for removing an electron.
The first ionization energy is the energy needed to remove one electron from the outermost shell.
The second ionization energy is the energy needed to remove a 2nd electron from the outermost shell after the 1st one has been removed.

18. Removing An Electron makes the atom an ion

19. Ionization Energy – Trends

20. Predicting Ionic Charges
Stable Noble gases do not form ions!
Group 18:

21. Predicting Ionic Charges
Group 1:
Lose 1 electron to form 1+ ions H+
Li+
Na+ K+

22. Predicting Ionic Charges
Group 2:
Loses 2 electrons to form 2+ ions
Be2+
Mg2+
Ca2+
Sr2+
Ba2+

23. Predicting Ionic Charges
Loses 3
electrons to form
3+ ions
Group 13:
B3+
Al3+
Ga3+

24. Predicting Ionic Charges
Lose 4
electrons or gain
4 electrons?
Group 14:
Neither! Group 13 elements rarely form ions.

25. Predicting Ionic Charges
Nitride
Gains 3
electrons to form
3- ions
Group 15:
P3-
Phosphide
As3-
Arsenide

26. Predicting Ionic Charges
Oxide
Gains 2
electrons to form
2- ions
Group 16:
S2-
Sulfide
Se2-
Selenide

27. Predicting Ionic Charges
F1-
Fluoride
Br1-
Bromide
Group 17:
Gains 1
electron to form
1- ions
Cl1-
Chloride
I1-
Iodide

28. The charge on an ion is also called its
Charge on Ions ? ?
The charge on an ion is also called its
Oxidation #

29. Ionic Radii Cations Anions Positively charged ions
Cations are smaller than their parent atoms because they lose electron(s) and also lose their outer energy level.
Anions
Negatively charged ions
Anions are larger than their parent atoms because they keep their outer energy level and they cannot hold their gained electron(s)
as tightly as their natural electrons

30. Unbalanced charges In normal atoms:
the # of Protons = the # of electrons
(positive and negative charges cancel each other out)
When an atom loses or gains an electron during chemical bonding the charges are unbalanced.
If an electron is lost, there is one fewer electron
# of protons > # of electrons
there is more positive charge than negative charge
The atom that donated the electron will now have one more proton than electrons is said to have a charge of plus 1 (+1).

31. Common Metals Cations

32. Common Nonmetals Anions

33. Atomic and Ionic radii Comparison

34. Ionization Energy - the energy required to remove an electron from an atom
Increases for successive electrons taken from
the same atom
Tends to increase across a period
Electrons in the same quantum level do
not shield as effectively as electrons in
inner levels
    Irregularities at half filled and filled
sublevels due to extra repulsion of
electrons paired in orbitals, making them
easier to remove
Tends to decrease down a group
Outer electrons are farther from the
nucleus

35. Ionization Energies

36. Think Greedy!!!!! Electronegativity
A measure of the ability of an atom in a chemical
compound to attract electrons
Electronegativities tend to increase across
a period
Electronegativities tend to decrease down a
group or remain the same
Think Greedy!!!!!

37. Electronegativity
Electronegativity is the ability of an atom to attract an electron
Shielding decreases the ability to attract an electron.
Cations do not want to attract electrons because they want,
instead, to lose electrons (the opposite thing). Therefore
cations have a low ability to attract an electron.
Anions easily attract electrons because they want to gain
electrons, making them high in electronegativity.
Noble gases, with their filled sublevels, do not want to attract
electrons so they have no ability to attract electrons and
therefore have no electronegativity.

38. Electronegativity

39. Shielding Effect The nuclear charge is shielded by the filled
inner electron
shells + + + +
The nucleus does not have enough
attraction to hold on to electrons
far away from the nucleus.

40. Atomic Radius
Slight Decrease
Increase

41. Ionization Energy
Increase
Decrease

42. Electronegativity
Increase
Decrease