1. The Periodic Table and Trends Topics 2 and 3
Please have a periodic table out.
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5. Dmitri Mendeleev 1834 – 1907 Early periodic tables…
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 numbers (Z) instead of mass
used x-ray spectroscopy to determine atomic # (# of protons)
also left blank spaces
properties now recurred periodically

9. Current periodic table

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

11. Know the names of these

12. 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

13. The highest
level is 4.

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

15. 2 1 3 5 8 2

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

17. Look, they are following my rule!

18. 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?

19. Periods 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 in the ratio when they react with other elements

20. IB loves the alkali metals and the halogens
Trends in the table

21. 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 electrostatic 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

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

23. 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. colors = ion
grey = neutral

25. periods across the periodic table
groups
increases downwards as more levels are added
more shielding
periods across the periodic table
radii decreases
the number of protons in the nucleus increases
this increases the strength of the positive nucleus and pulls electrons in every 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

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

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

31. anions (- ions) are larger than parent atom
have gained an electron to achieve a stable 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. trends once they are ions across a period
cations decreases at first just like atoms (stronger nucleus pulling in levels) AND ratio of protons to electrons increases
suddenly increases when become anions
then goes back to decreasing just like atoms (stronger nucleus pulling in levels) AND ratio of protons to electrons increases
down a group (same as neutral atoms)
increases as new levels are added
more levels of shielding

33. dark grey is the size of the ion

34. colors = ion
grey = neutral

35. ionic trend across period 2

36. Ionization energy
the minimum energy (kJ mol-1) 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…

37. don’t forget-- gaseous

38. 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

40. increases across a period
the nucleus is becoming stronger (effective nuclear charge) and therefore all levels of electrons are pulled closer
atomic radii is decreasing
requires 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

42. the increase in ionization energy is not continuous across the table
electrons are also harder to remove…
a sub-level (s,p,d,f) is completely filled
a sub-level (s,p,d,f) are half filled

43. notice discontinuity as move across period 3 only last sub-level shown
full sub-level
notice discontinuity as move across period 3
½ full sub-level
full sub-level
only last sub-level shown

44. 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.

46. 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

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

49. Electron Affinity (Ea)
the change in energy (kJ/mole) of a neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative ion (anion)
in other words, the neutral atom's likelihood of gaining an electron
example
F(g) + e-  F-(g) will release 328 kJ/mole of energy (therefore – 328 kJ/mol)
the more negative Ea, the greater the attraction for the electron

50. trends across a period down a group
in general, Ea becomes more negative from L to R
all the same reasons as before (effective nuclear charge, distance…)
down a group
in general, becomes less negative
all the same reasons as before (shielding, effective nuclear charge, distance…)

52. the more negative, the more energy released when an electron is added
positive means it would need energy to accept an electron

53. next trends require understanding of concepts covered in later topics (this year and even senior year)
only need to know the trends for the unit test, NOT the reason why until later in the year

54. Melting point down group 1 (alkali metals) down group 17 (halogens)
decreases as “sea of negative
electrons” are farther away
from the positive metal ions
down group 17 (halogens)
increases as the van der Waals’ forces increase
larger molecules have more
electrons which increases
the chance that one side of
the molecule could be negative
Element
Melting Point (K) Li
453 Na
370 K
336 Rb
312 Cs
301 Fr
295

56. across the table (period 3)
from left to right
increases until group 14 (C is in group 14, think diamonds) then decreases starting at group 15
bonding goes from strong metallic to very strong macromolecules (giant network covalent) to weak van der Waals’ attraction

57. = increases

58. Chemical properties groups alkali metals
react vigorously with water and air
2Na (s) + H2O (l)  2Na (aq) + 2OH- (aq) + H2 (g)
(Li, Na, K… all the same equation)
reactivity increases downwards
because the outer (valence) electron is in higher energy levels (farther from the nucleus) and easier to remove
react with the halogens
halogens’ reactivity increases upwards
smaller size attracts electrons better
since they can be close to the
nucleus
1+ charge
1- charge

59. least reactive
most reactive

60. None of the following slides are on the test.
They will be covered later in the year and next year.

61. halogens (group 7) diatomic molecules such as Cl2, Br2, I2
can react with halide ions (Cl -, Br -, and I -)
the most reactive ends up as an ion (1- charge) and is not visible (molecules Cl2, Br2, I2 are a visible gas)
Cl > Br > I
Cl-(aq)
Br-(aq)
I-(aq)
Cl2
Colorless- no reaction
turns red due to formation of Br2
turns brown due to formation of I2
Br2
no reaction I2

62. periods from left to right in period 3 metals…metaloids…nonmetals
oxides are
ionic…..and then covalent bonds
when oxides react with water
basic…amphoteric (either basic or acidic)…acidic
Na2O(s) + H2O (l)  2 NaOH (aq) strong base
MgO (s) +H2O (l)  Mg(OH)2 (aq) weaker base
P4O10 (s) + 6H2O (l)  4 H3PO4 (aq) weak/strong acid
SO3(g) + H2O (l)  H2SO4 (aq) strong acid

63. Look at the blue arrows! Senior year…