1. Periodic Properties of the Elements

2. The Periodic Table
The modern periodic table was developed in 1872 by Dmitri Mendeleev ( ). A similar table was also developed independently by Julius Meyer ( ).
The table groups elements with similar properties (both physical and chemical) in vertical columns. As a result, certain properties recur periodically.

3. The Periodic Table
Mendeleev left empty spaces in his table for elements that hadn’t yet been discovered. Based on the principle of recurring properties, he was able to predict the density, atomic mass, melting or boiling points and formulas of compounds for several “missing” elements.

4. The Periodic Table

5. The Periodic Table
metal/non-metal line

6. The Periodic Table
The periodic table is based on observations of chemical and physical behavior of the elements. It was developed before the discovery of subatomic particles or knowledge of the structure of atoms. The basis of the periodic table can be explained by quantum theory and the electronic structure of atoms.

7. Periodic Trends
Many of the properties of atoms show clear trends in going across a period (from left to right) or down a group.
In going across a period, each atom gains a proton in the nucleus as well as a valence electron.

8. Periodic Trends
The increase of positive charge in the nucleus isn’t completely cancelled out by the addition of the electron.
Electrons added to the valence shell don’t shield each other very much. As a result, in going across a period, the effective nuclear charge (Zeff) increases.

9. Effective Nuclear Charge
The effective nuclear charge (Zeff) equals the atomic number (Z) minus the shielding factor (σ).
Zeff= Z-σ

10. Effective Nuclear Charge
The effective nuclear charge (Zeff) equals the atomic number (Z) minus the shielding factor (σ).
Zeff= Z-σ
Within the valence shell, the shielding factor is approximately 0.35, so going across a period results in an increase in Zeff of roughly .65 for each element.

11. Effective Nuclear Charge
Zeff= Z-σ

12. Effective Nuclear Charge
Electrons in the valence shell are partially shielded from the nucleus by core electrons.

13. Effective Nuclear Charge
Electrons in p or d orbitals don’t get too close to the nucleus, so they are less shielding than electrons in s orbitals. As a result, effective nuclear charge increases across a period.

14. Periodic Trends

15. Periodic Trends
In going down a group or family, a full quantum level of electrons, along with an equal number of protons, is added.
As n increases, the valence electrons are, on average, farther from the nucleus, and experience less nuclear pull due to the shielding by the “core” electrons. As a result, Zeff decreases slightly going down a group.

16. Trends- Atomic Radii Atomic radii are obtained in a variety of ways:
1. For metallic elements, the radius is half the internuclear distance in the crystal, which is obtained from X-ray data.
2. For diatomic molecules, the radius is half the bond length.
3. For other elements, estimates of the radii are made.

17. Trends- Atomic Radii
Atomic radii follow trends directly related to the effective nuclear charge. As Zeff increases across a period, the electrons are pulled closer to the nucleus, and atomic radii decrease.
As Zeff decreases down a group, the valence electrons experience less nuclear attraction, and the radius increases.

18. Trends- Atomic Radii
Atomic size roughly halves across a period, and doubles going down a group.

19. Trends – Ionization Energy
Ionization energy is the energy required to remove an electron from a mole of gaseous atoms or ions.
X(g) + energy  X+(g) + e-
Elements can lose more than one electron, so there are 1st, 2nd, 3rd, etc., ionization energies.

20. Ionization Energy
It always requires energy to remove an electron from a neutral atom.
As more electrons are removed and the ion becomes positively charged, it requires increasingly greater energy to remove electrons.

21. Trends – Ionization Energy
Ionization energy is a measure of how tightly the electrons in the highest occupied orbitals are held by the nucleus. As a result, it is directly related to the effective nuclear charge.
Ionization energy increases going across a period, and decreases going down a group.

22. Trends – Ionization Energy

23. Trends – Ionization Energy

24. Ionization Energy

25. Trends – Electron Affinity
Electron Affinity involves the addition of an electron to a mole of gaseous atoms. There are different conventions to defining electron affinity. Your text defines the EA as the energy released during the following process:
X(g) + e-  X-(g)

26. Electron Affinity
Your text defines the EA as the energy released during the following process:
X(g) + e-  X-(g)
A positive value for EA indicates that the process releases significant energy. Thus, the halogens tend to have high electron affinities.

27. Trends – Electron Affinity
There is less of a predictable trend in electron affinities. In going across a period (ignoring the noble gases), the electron affinity should become more negative. Although this is observed, there are many inconsistencies.

28. Trends – Electron Affinity

29. Trends - Electron Affinity

30. Trends- Electron Affinity
In going down a group, the electron affinity should become become smaller. Although this trend is observed, there is only a slight change in electron affinities within a group. There may also be inconsistencies in the general trend.

31. Electron Affinity
It should be noted that the addition of a second electron to an anion is always highly unfavorable. The electron affinity of oxygen is 141 kJ/mol to form O–. Addition of the second electron to form the oxide ion (O2–) requires 744 kJ/mol.

32. Metallic Character
Metals are shiny, malleable and ductile. They are generally good conductors of heat and electricity, and low ionization energies. In reaction with non-metals, metals tend to lose electrons and form cations.

33. Metallic Character

34. Metallic Character
Across a period, metallic behavior decreases. Non-metals are often crumbly solids, liquids or gases at room temperature.

35. Metallic Character
Metallic behavior increases going down a group.

36. Electron Configurations of Ions
The atoms of the main group elements (groups IA-VIIA) will form ions by losing or gaining electrons. The resulting ion will have the same electron configuration as a noble gas (group VIIIA). These configurations are usually very stable.

37. Electron Configurations of Ions
Atoms or ions with the same electron configuration (or number of electrons) are called isoelectronic.
For example, Na+, Mg2+, Ne, F-, and O2- are isoelectronic. The size will decrease with increasing positive charge.
O2- > F- >Ne> Na+> Mg2+

38. Electron Configurations of Ions
When atoms lose electrons, the electrons are always removed from the highest quantum level first.
For the first row of transition metals, this means that the electrons in 4s subshell are lost before the 3d subshell.
Fe: [Ar]4s23d6 Fe2+: [Ar] 3d6 or [Ar]4s03d6

39. Common Ionic Charges
The charges of ions of elements in groups 1A-7A (the main groups) are usually predictable.
Group 1A metals form +1 ions, group 2A metals form +2 ions, etc.
The non-metals of group 5A have a -3 charge, those of group 6A have a -2 charge, and the halogens form ions with a -1 charge.

40. Typical Ionic Charges

41. Trends – Ionic Size
Cations are always smaller than the neutral atom. The loss of one or more electrons significantly increases Zeff, resulting in the valence electrons being pulled closer to the nucleus.

42. Ionic Size - Cations
Within a group, assuming the same ionic charge, the size of the ion increases going down the group, due to more core electrons shielding the nucleus as n increases.

44. Trends – Ionic Size
Across period, the cations get more positive, and as a result, considerably smaller.

45. Trends – Ionic Size
Anions are always larger in size than the neutral atom. The addition of one or more electrons results in greater electron-electron repulsion, which causes the valence electrons to “spread out” a bit.

46. Size of Anions

47. Anions are always larger than the neutral atom.

48. Size of Anions
Within a group, assuming the same ionic charge, the size of the ion increases going down the group, due to more core electrons shielding the nucleus as n increases.

49. Trends – Ionic Size

50. Group IA – the Alkali Metals
In discussing the chemistry, preparation and properties of the group IA elements, it is important to remember that hydrogen is not a group IA metal. It’s properties and reactivity would place it within group 7A (diatomic non-metals), rather than group IA.

51. Group 1A Metals
The group 1A metals are soft shiny metals with fairly low densities (Li, Na and K are less dense than water) and low melting points. Sodium melts at 98oC, and cesium melts at 29oC.
The softness, low density and low melting points are the result of weaker metallic bonding due to only one valence electron in this group.

52. Group 1A Metals - Production
Due to the high reactivity with oxygen and water, all of the metals are found in nature in ionic form (M1+).
The pure metal must be produced in an oxygen and water-free environment. Typically, an electrical current is passed through the melted chloride salt. The metal and the chlorine gas are collected separately.

53. Reactivity Trends
The chemical behavior of the group IA metals illustrates periodic trends. As the valence electron occupies a higher quantum level, it experiences less nuclear attraction, and is more easily removed.

54. 2 M(s) + 2 H2O(l)  H2(g) + 2 MOH(aq)
Group 1A Metals + Water
The reaction with water forms hydrogen gas and the aqueous metal hydroxide. The reaction is so vigorous, that the hydrogen may ignite.
2 M(s) H2O(l)  H2(g) MOH(aq)

55. Metallic Character
The group IA metals react with water to produce hydrogen and the metal hydroxide.
Metallic behavior increases going down a group.