1. Periodic Properties of the Elements
Ch 7

2. History of the Periodic Table
Johann Dobereiner (early 1800’s) designed the first chart showing triads. 3 elements with similar properties.
John Newland (1864) suggested that the properties repeat every 8th element, Octaves.
Meyer & Dmitri Mendeleev (1869) designed the periodic table bases on the atom mass.
Henry Mosley (early 1900’s) designed the modern periodic table bases on the atom number.

3. Early 1900’s Henry Mosley organized the elements according to their atomic number (number of protons). This is the modern periodic table that we use.

5. Remember Coulomb’s Law: the strength of the interaction between 2 electrical charges depends on the magnitude of the charges and the distance between them.
The force increases as the nuclear charge increases
The force decreases as the electron moves farther from the nucleus

6. Effective Nuclear Charge (Shielding effect) – the electron experiences a net attraction that is the result of the nuclear attraction decreased by the electron – electron repulsion.
Zeff = Z – S
Effective nuclear charge = protons – core electrons
Na has 11 p+, 11 e-
Only 1 e- is in the outer energy level, 10 inner core e-
Zeff = 11 – 10
Zeff = 1
The 3s e- experiences only the pull of 1 p+
Zeff increases  on Periodic Table,
and increases as go down the PT

7. Atomic radius: Size of the radius decreases as you go from left to right on the periodic table. Due to an increased nuclear charge, Effective nuclear charge or shielding effect. Size of the radius increases as you go down the family There is little change in the size as you go through the transition elements. *Cation (+ ions) become smaller than the atomic size for that element. *Anion (- ions) become larger than the atomic size for that element.

8. Nonbonding atomic radius or Van der Waals radius.

10. Cation & Anion sizes

11. Isoelectric Series Cation : since they lose e-, they tend to be smaller than the parent atom. Anion: since they gain e-, they are significantly larger than the parent atom. Isoelectric ions: ion with the same number of e-
Example: O-2 , F-1 , Na+1 , Mg+2 , Al+3 These all have the same number of e- as Ne e- repulsion is the same in all! As the number of p+ increases, the size of the ion decreases. O-2 F-1 Na+1 Mg+2 Al+3
140 pm 136 pm 95 pm 65 pm 50 pm
All having 10 e-

12. Periodic Trends
See page 270
Ionization energy
Ionizations Energy: The energy required to remove an e- from a gaseous atom or ion.
X(g)  X e- DH = (+)
Alkali metals have the lowest ionization energy
Noble gases have the highest ionization energy
Progressive ionizations (removal of more than 1 e-) increase in energy. There is a drastic increase when trying to remove an e- from a full energy level.

15. Notice that removing e- from a partially filled energy level is a relatively low energy amount. As more e- are removed the energy increases. To remove an e- from a full lower energy level results in a large jump in the energy required.

16. Electron Affinity is the energy associated with the addition of an e- to a gaseous atom. X(g) + e-  X-1 DH = (-) e- affinity becomes more exothermic (- value) as you move from the left to the right across the periodic table. * see page 272

17. e- affinity

18. Classification of elements as metals, nonmetals, and matalloids.

19. Characteristic Properties
Metals
Nonmetals
Shiny silvery solids Malleable Ductile Conduct heat & electricity Metal oxide + H2O  basic Forms cations
Various colors & phases Solids usually brittle some hard, some soft Poor conductor of heat & electricity Nonmetal oxide + H2O  acidic Forms anions

21. Diatomic molecules
There are 7 elements that do not exist as 1 atom alone. As an element, they are always as 2 atoms bonded together as a molecule.
Location of the
diatomic elements

22. Alkali Metals 2 M(s) + 2 H2O(l)  2 MOH(aq) + H2(g)
All characteristic metal properties (soft solids)
Combine directly with nonmetals
2 M(s) + H2 (g)  2 MH(s)
2 M(s) + S(s)  M2S(s)
React vigorously with H2O to release H2
2 M(s) H2O(l)  2 MOH(aq) + H2(g)
React with O2 to form oxides
4 M(s) + O2(g)  2 M2O(s)

23. Flame Testing to ID alkali metals
Lithium Sodium Potassium
Crimson red Yellow Lilac
Flame Testing to ID alkali metals

24. Chart showing the similarities of properties within the Alkali Metal family.

25. Alkali Earth Metals All characteristic metal properties
React with H2O to release H2(g) (less vigorous)
M(s) H2O(l)  M(OH)2(aq) + H2(g)
React with halogens to form salts
M(s) + Cl2(g)  MCl2(s)
React with O2 to form oxides
2 M(s) + O2(g)  2 MO(s)
Flame test ID: Sr – brilliant red
Ba - green

26. Chart showing the similarities of properties within the Alkali Earth Metal family.

27. Hydrogen HCl(g) + H2O(l)  H3O+1(aq) + Cl-1(aq)
Colorless diatomic gas
Tends to react with nonmetals sharing e-
H2(g) + Cl2(g)  2 HCl(g)
Readily forms H+1 in the presence of H2O to form acids
HCl(g) + H2O(l)  H3O+1(aq) + Cl-1(aq)
Reacts with metals to form hydrides
2 M(s) + H2 (g)  2 MH(s)

28. Oxygen group
O, S, Se typical nonmetals, Te is metalloid, and Po radioactive metal
O2 colorless gas, others solid
O has allotropes, different forms of same element: O2 and O3
O3 has pungent odor, strong oxidizing agent (great tendency to attract e-)

29. Chart showing the similarities of properties within the Oxygen family.

30. Halogens Typical nonmetals : F, Cl are gases, Br a liquid, I a solid
All are diatomic
Tendency to gain e- to form halide ions
SiO2(s) + 2F2(g)  SiF4(g) + O2(g)
React with metals to form ionic halides
2 M(s) + H2 (g)  2 MH(s)

31. Diatomic elements
Iodine (I2) purple
Bromine (Br2) orange
Chlorine (Cl2) colorless

32. Chart showing the similarities of properties within the Halogen family.

33. Noble Gases
Monatomic gases
Inert

34. Chart showing the similarities of properties within the Noble gases.

35. Magnetic Properties
Diamagnetic elements have all e- in pairs, therefore all individual magnetic properties resulting from electrical spin is cancelled.
Diamagnetic elements are repelled by magnetic fields so they will weigh slightly less in a magnetic field.
Paramagnetic elements have unpaired e-, therefore the individual magnetic properties resulting from electrical spin is not cancelled. There is an induced magnetic field due to the spin of the unpaired e-.
Paramagnetic elements are attracted by a magnetic field so they will weigh slightly more in a magnetic field.

36. Common Ions