1. Chemical Periodicity

2. Mendeleev’s Periodic Table
In 1871, Dmitri Mendeleev arranged all of the then-known elements by order of increasing atomic mass.

3. The Modern Periodic Table
In 1913, Henry Moseley created a more accurate positioning of the elements in the modern periodic table. He arranged the elements in order of increasing atomic number.
He was killed in action at Gallipoli in 1915 at the age of 28.

4. The Modern Periodic Table

5. The Periodic Table Remember… Vertical columns are groups.
Horizontal rows are periods.
and…
Valence electrons – electrons in the outermost (highest) principal energy level of an atom
Core electrons – inner electrons

6. Patterns in the electron configurations of the elements
The noble gases:
these are elements in which the outermost s and p sublevels are completely filled.
He: 1s2
Ne: 1s2 2s2 2p6
Ar: 1s2 2s2 2p6 3s2 3p6

7. Patterns in the electron configurations of the elements
The representative elements:
these elements have the outermost s and p sublevel only partially filled.
Li: 1s2 2s1
Na: 1s2 2s2 2p6 3s1
K: 1s2 2s2 2p6 3s2 3p6 4s1

8. Patterns in the electron configurations of the elements
The transition metals:
In these metallic elements, the outermost s and nearby d sublevel contain electrons.
The inner transition metals:
In these metallic elements, the outermost s and nearby f sublevel generally contain electrons.

9. Patterns in the electron configurations of the elements
Elements with the same valence electron arrangement show very similar chemical behavior.

10. Electron Configurations and the Periodic Table
Electron configurations for K through Kr

12. Practice writing electron configurations from the periodic table.

13. Covalent atomic radius
It is not possible to directly measure the size of an atom.
Chemists determine the distance between the nuclei in atoms of diatomic elements, then divide this by 2 to find the radius of one atom.

14. I. Periodic Trend in Atomic Size
Atomic radius tends to increase down the group.
Atomic radius tends to decrease across the period.

15. Atomic radius

16. Atomic radius

17. Reason for the trends
Atomic size decreases across the period because the effective nuclear charge increases across the period.
Atomic size increases down the group because the shielding effect increases down the group.
(Shielding effect occurs as the core electrons shield the attraction of the nucleus on the outermost electrons.)

18. II. Periodic Trend in Ionization Energy (IE)
Ionization Energy (IE) is the energy required to remove an electron from a gaseous atom.
Na(g)  Na+(g) + e-
Ionization energy

19. Ionization Energy (IE) in kilojoules/mole
First IE is the energy required to remove the first electron.
Second IE is the energy required to remove the second electron (NOT both first + second), etc.

20. Explain why there is a large increase between the first and second IE for Na.
Explain why there is a large increase between the second and third IE for Mg.
Predict where there will be a large increase in IE for Al.

21. Ionization Energy (IE)
Ionization energy tends to decrease down a group.
Ionization energy tends to increase across a period.
Explain why the trend is as described (same reason as for the size trend).

22. Explain why these elements do not follow the trend...
Consider the electron configurations as you explain why the trend is not always consistent as you move across the period (such as across period 2, from Li to Ne).
Why is the IE lower than predicted by the trend for B? …for O?

23. Ionization energy

24. IE in Metals
Metals tend to lose electrons to form positive ions because their ionization energies are low.

25. III. Periodic Trend in Electron Affinity (EA)
Electron affinity (EA) is the energy change that accompanies the addition of an electron to a gaseous atom.
Br(g) + e-  Br-(g)
Electron affinity

26. Electron Affinity
Electron Affinities for the Representative Elements (in kilojoules/mole)

27. Electron Affinity Electron Affinity tends to decrease down a group.
Electron Affinity tends to increase across a period.
Explain why the trend is as described (same reason as for the size trend).

28. Electron affinity

29. IV. Periodic Trend in Ionic Size
Cations are smaller than their original atoms. (Remember; atoms lose electrons to form “+” cations.)
Anions are larger than their original atoms. (Atoms gain electrons to form “-” anions)
Ionic radius increases down the group (as does atomic radius).

30. Ionic Size

31. V. Periodic Trend in Electronegativity
Electronegativity – the tendency for the atom to attract electrons to itself in a chemical bond
Increases from left to right across a period
Decreases down a group

32. Pauli Electronegativity Values

33. Electronegativity

34. Electronegativity
Linus Pauling established the concept of electronegativity and developed a scale that helped to predict the nature of chemical bonding.
In 1954, Pauling was awarded the Nobel Prize in Chemistry.

35. Summary of Periodic Trends

36. Review Group Names Remember: Group 1A = Alkali metals
Group 2A = Alkaline earth metals
Group 7 A = Halogens
Group 0 = Noble Gases