1. Trends on the Periodic Table

2. Periodic Trends
The arrangement of the periodic table reveals trends or general tendancies in the physical and chemical properties of the elements.
Periodic Law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers.
Periodic means REPEATING.

3. What is a Trend?
A trend is a predictable change in a particular direction.
For example, as you move down Group 1, reactivity increases for each element.
Knowing these various trends allows you to make logical predictions about the elements chemical and physical properties.

4. Periodic Trends in Atomic Radii
The exact size of an atom is difficult to determine.
The size is determined by the volume occupied by the electrons surrounding the nucleus.
We model atoms as clouds, but in reality there is no clear cut boundaries for the edges of atoms.

5. Radii of Atoms
Are usually determined for atoms that are chemically bonded or close together in a solid state.
Bond radius is half the distance from the center of one like atom to the center of another like atom bonded chemically together.
Bond radius is used to determine the size of an atom.

6. Bond Radius Equation

7. Another Way to Determine Size of an Atom
Van der Waals is half the distance between the nuclei in adjacent non-bonded molecules. Sometimes this radius is used to denote the size of atoms, but is available only for the main group elements.

8. Van der Waals Forces

9. The Problem with Bond Radius is…
That bonds between different atoms create different size radii.
For example, bonds between Tin atoms in metallic Tin are different than Tin Chloride, SnCl4. But even with this limitation, it is still a useful way to measure the size of atoms.

10. Atomic Radii Trends Atomic Radii increases down a group.
There is a general trend toward LARGER RADII as you proceed down a group.

11. Atomic Radii on the Periodic Table

12. Reason for Larger Radii
1. The addition of one more principle energy level from one period to the next.
2. The phenomenon of shielding.
Shielding-is the reduction of the attractive force between a positively charged nucleus and its outermost electrons due to the cancellation of some of the positive charge by the negative charge of the inner electrons.

13. Shielding
Electrons in the inner energy levels “shield” the outer electrons from the attractive force of the positive nucleus. Hence, these inner electrons reduce the positive force of the nucleus so the outer electrons can move even further out from the nucleus.

14. Shielding

15. Conclusion
So the outermost electrons (valence electrons) are not subject to a full charge of the nucleus, they, therefore are not held as close to the nucleus.

16. Nuclear Charge
As you move down a group, nuclear charge (positive charge) increases because more protons are present.
Also, at the same time, more shielding is present do to the presence of more electrons.
This creates a canceling effect regardless of energ level.
presence of more inner electrons.
This creates a canceling effect regardless of the energy level.
This means shielding is constant all the way down a group.

20. Example
Sodium, Na and Cesium, Cs have about the same net nuclear charge even though Cesium has more protons in the nucleus. Cesium also has more electrons to create a greater shielding effect.

21. Atomic Radii Decrease across a Period
Reason for Smaller Radii
1. It is caused by increasing positive charge of the nucleus-increased nuclear charge. You gain one more proton than the element before.
2. You have added electrons however they are going in the same energy level. So therefore, shielding remains constant There is no cancellation of the positive charge like there is when you move down a group.

22. Conclusion
So therefore, more protons are added without a canceling effect from the electrons. Since no more energy levels are being added to distance the electrons from the positive nucleus, the electrons in the outer energy level get pulled in closer.

23. This Trend Levels Off
This trend gets less pronounced when there are many electrons between the nucleus and the outermost energy level. When many electrons are pulled in closer to the nucleus, the electrons get too close to each other and start to repel each other.
See Period Six

25. Conclusion
Therefore, there is a point where the electrons will come no closer to each other and they will level off in size.

26. Ionization Energy & Electron Affinity
Ionization Energy is the energy you use to remove an outer electron (valence electron) from an atom.
Equation:
A IE -> A e-
Atom + ionization energy yields cation + electron

27. Ionization Energy

29. Trends in Ionization Energy
Ionization Energy increases across a Period.
Reason:
1. More protons are added to the nucleus as you move across a period, but they are added to the same energy level.
2. The nuclear charge (positive charge) increases holding the electrons closer because shielding is constant.

30. Conclusion
Therefore, you need MORE IONIZATION ENERGY to pull electrons off the atom!

31. Ionization energy Decreases down a Group
Reason:
1. The number of energy levels increase between the nucleus and outermost electrons.
2. The outer shell electrons are further from the nucleus.
3. The net nuclear charge (positive charge) is the same all the way down the group due to increased shielding.

32. Conclusion
Therefore, less energy is required to remove an electron. The attractive forces between the nucleus and the outermost electrons decrease and the energy necessary to remove them also decreases.

33. Periodic Trends in Electron Affinity
Electron Affinity is the energy emitted (given off) upon the addition of an electron to an atom or group of atoms in the gas phase.
Because inner electrons do not shield the positive nucleus one hundred percent, an approaching electron may experience a net pull and be drawn into a vacant orbital in an atom.

34. Electron Affinity

35. Equation A + e- -> A- + energy
Atom + electron yields anion + energy given off

36. Electron Affinity Increases going across a Period
Reason:
1. Shielding is constant because the energy level is the same.
2. Nuclear charge (positive charge) increases due to the addition of one more proton.

37. Conclusion
Therefore, the atom’s attraction for extra electrons increases left to right.

38. Electron Affinity Decreases down a Group
Reason:
1. Both shielding and nuclear charge increase, but each cancel the other making net nuclear charge constant.
2. This constant nuclear charge allows electrons to move further out from the nucleus.

39. Conclusion
The atom’s attraction for extra electrons decreases!

40. Electronegativity of an Atom
Is the tendancy of an atom to attract electrons to itself when it is chemically combined with another element.

41. Electronegativity is expressed…
In terms of a relative scale with arbitrarily selected standard units.
Fluorine is the most electronegative element and is assigned 4.0 as the largest unit.
Francium is the least electronegative and is assigned a value of 0.7.

42. Trends in Electronegativity
Electronegativity increases across a period.
Reason:
1. Nuclear charge increases as you move across a period. This is due to the fact that shielding remains constant.

43. Electronegativity Chart

45. Conclusion
Therefore, the atom attracts electrons to itself the greater the nuclear charge.

46. Electronegativity Decreases down a Group
Reason:
1. Because increased shielding is present, nuclear charge remains relatively constant down a group.

47. Conclusions
Therefore, atoms further away from the nucleus are less likely to be attracted by the nuclear charge and the atom’s attraction for more electrons decreases!
Special Note: Chemists determine electronegativity values by measuring the polarities of the bonds between various atoms.