1. The Periodic Table and Trends
Please have a periodic table out.
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2. Use this one.

5. Dmitri Mendeleev 1834 – 1907 Russian chemist and teacher
given the elements he knew about, he organized a “Periodic Table” based on increasing atomic mass (it’s now atomic #)
he even left empty spaces to be filled in later

6. At the time the elements gallium and germanium were not known
At the time the elements gallium and germanium were not known. These are the blank spaces in his periodic table. He predicted their discovery and estimated their properties.

7. Henry Moseley – 1915
arranged the elements in increasing atomic number (Z)
properties now recurred periodically better than before

9. Design of the Table Groups (18 total) are the vertical columns.
elements have similar, but not identical, properties
most important property is that
they have the same # of valence
electrons

10. Know the names of these groups

11. valence electrons- electrons in the highest occupied energy level
all elements have 1,2,3,4,5,6,7, or 8 valence electrons
These are a lower level. Therefore the d sub-level is never included for valence electrons

12. The highest
level is 4.

13. Lewis Dot-Diagrams/Structures
a short cut wherevalence electrons are represented as dots around the chemical symbol for the element Na Cl

14. 2 1 3 5 8 2

15. What two blocks will always be the highest occupied level?

16. Look, they are following my rule!

17. Br is [Ar] 4s2 3d10 4p5 How many valence electrons are present?
B is 1s2 2s2 2p1;
2 is the outermost energy level
it contains 3 valence electrons, 2 in the 2s and 1 in the 2p
Br is [Ar] 4s2 3d10 4p5 How many valence electrons are present?

18. Periods (7) are the horizontal rows
do NOT have similar properties
however, there is a pattern to their properties as you move across the table that is visible when they react with other elements

19. Trends in the table

20. many trends are easier to understand if you comprehend the following
the ability of a nucleus to “hang on to” or attract its valence electrons is the result of two opposing forces
the attraction between the electron (-) and the nucleus (+)
the repulsions between the electron (-) in question and all the other electrons (-) in the atom (this is called the shielding effect)
the net resulting force of these two is referred to effective nuclear charge

21. This is a simple, yet very good picture. Do you understand it?

22. Atomic radii the distance from the nucleus to the outermost electron
cannot measure the same way as a simple circle due to electrons are not in a fixed location
therefore measure distance between two nuclei and divide by two

24. periods across the periodic table
groups
increases downwards as more levels are added
more shielding pushes outer levels out
periods across the periodic table
radii decreases
the number of protons in the nucleus increases
increases the strength of the positive nucleus and pulls electrons in the given level closer to it
added electrons are not contributing to the shielding effect because they are still in the same level H Li Na K Rb
McGraw Hill video

28. Looking at ions compared to their parent atoms
Ionic radii
Looking at ions compared to their parent atoms
atoms tend to gain or loose electrons in order to have the electron configuration of a noble gas

29. cations (+ ions) are smaller than the parent atom
have lost an electron (actually, lost an entire level!)
therefore have fewer electrons than protons
less no shielding
Li+ .078nm +
Li forming a cation Li
0.152 nm

30. anions (- ions) are larger than parent atom
have gained an electron(s) to achieve noble gas configuration
effective nuclear charge has decreased since same nucleus now holding on to more electrons
plus, the added electron repels the existing electrons farther apart (kind of “puffs it out”) F -
0.133 nm
10 e- and 9 p+
F nm
9e- and 9p+

32. Ionization energy
the minimum energy (kJ/mol) needed to remove an electron from a neutral gaseous atom in its ground state, leaving behind a gaseous ion
X(g)  X+(g) + e-
first ionization energy- energy to remove first electron
second ionization energy- energy to remove second electron
third ionization energy- and so on…

33. don’t forget-- gaseous

34. decreases down a group
outer electrons are farther from the nucleus and therefore easier to remove
inner core electrons “shield” the valence electrons from the pull of the positive nucleus and therefore easier to remove

36. increases across a period
the nucleus is becoming stronger (effective nuclear charge) and therefore the levels, especially valence electrons, are pulled closer
atomic radii is decreasing
this makes it harder to remove a valence electron since it is closer to the nucleus
or another way to look at it… a stronger nuclear charge acting on more contracted orbitals

38. Electronegativity
measures the attraction for a shared pair of electrons in a bond
Linus Pauling (1901 to 1994) came up with a
scale where a value of 4.0 is arbitrarily
given to the most electronegative element,
fluorine, and the other electronegativities
are scaled relative to this value.

40. trends (same as ionization energy and for the same reasons)
as you go down a group electronegativity decreases
the size of the atom increases
the bonding pair of electrons (-) is increasingly distant from the attraction of the nucleus (+)
the bonding pair of electrons (-) are shielded because of core electrons (-) interfering with the nucleus’ (+) hold on valence electrons H Li Na K Rb

42. as you go across a period
electronegativity increases
the atoms become smaller as the effective nuclear charge increases
easier to attract the shared pair of electrons as they will be in an orbital closer to the nucleus moving from L to R on the table