1. Trends in the Periodic Table

2. History of the Periodic Table
The first Periodic Table was devised by Dmitri Mendeleev
Table was based on the number of protons
Some elements were missing (under Si)
Predicted properties based on trends that were observed
1886 – Germanium was discovered – predicted properties of Ge matched experimental properties

3. Introduction to Trends
Elements are arranged to show similarities and differences in properties
Properties are related to:
Number of protons (Nuclear Charge)
Number of energy levels

4. Introduction to Trends
Nuclear charge: the nucleus contains positively charged protons, which attract negatively charged electrons. As the atomic number increases, the number of protons increases and therefore the nuclear charge increases. The higher the nuclear charge, the more strongly “held” (attracted) the valence shell electrons are to the nucleus.
+++
+++
Lower nuclear charge
weaker attraction
Higher nuclear charge
stronger attraction

5. Introduction to Trends
Energy levels (shells): Electrons are located outside of the nucleus in specific energy levels. These energy levels are a certain distance from the nucleus. The further from the nucleus the electrons are located, the less tightly “held” (attracted) to the nucleus they are.
Row Number = Period Number = Numbers of Energy Levels
+++
Less energy levels
Valence electron closer to the nucleus
Less shielding
Stronger attraction
More energy levels
Valence electron further from the nucleus
More shielding
Weaker attraction
+++

6. Atomic Radius
As you go across the table (from left to right), the size of the atom decreases
Reason: Higher nuclear charge “pulls” more on the outer valence shell electrons, making the overall size decrease
3 p p p p p p p p+

7. Atomic Radius Going down the table, the size increases
Reason: More energy levels
weakens attraction between
valence shell electrons and
nucleus, making the overall
size increase

8. Question: Arrange the following atoms in order of
increasing atomic radius (from smallest to
largest): Sc, Ba, Se.

9. Answer: Se < Sc < Ba 4 shells 4 shells 6 shells 34 p+ 21 p+
less shielding  stronger more shielding  weaker “pull”
“pull” on valence shell e on valence shell e-
higher nuclear
charge  stronger
pull on e-

10. Atomic Radius
decreases
increases

11. Ionization Energy (I.E.)
Amount of energy required to remove ONE electron from an atom (in a gaseous state)
+++
+++
ACROSS THE
PERIODIC TABLE
Lower nuclear charge,
weaker attraction
Higher nuclear charge, stronger attraction
Easier to remove e-
More difficult to remove e-
Lower I.E.
Higher I.E.

12. Ionization Energy (I.E.)
Amount of energy required to remove ONE electron from an atom (in a gaseous state)
Less energy levels
Valence electron closer to the nucleus
Less shielding
Stronger attraction
Higher I.E.
+++
DOWN THE
PERIODIC TABLE
More energy levels
Valence electron further from the nucleus
More shielding
Weaker attraction
Lower I.E.
+++

13. Atomic Radius
I. E.
decreases
increases
increases
decreases

14. Question: Arrange the following in order of increasing
ionization energy (lowest I.E. to highest I.E.):
As, Sn, Br, Sr.
Explain your reasoning.

15. Answer: Sr < Sn < As < Br 5 shells 5 shells 4 shells 4 shells
38 p p p p +
more shielding  less less shielding  stronger “pull”
“pull” on valence shell e on valence shell e-
higher nuclear
charge  stronger
pull on e-

16. Atomic Radius I. E. decreases increases (non-metals: don’t want to
lose e-, therefore high I.E.)
increases
decreases
(metals: want to lose e-,
therefore low I.E.)

17. Electron Affinity (E. A.)
A change in energy that accompanies the addition of an electron to an atom in the gaseous state
High affinity : High attraction; Easy to gain electron
Low affinity: Low attraction; Difficult to gain electron

18. Atomic Radius I. E. E. A. decreases increases increases non-metals
want to gain
e-, therefore
high E.A.
increases
decreases
decreases
metals
do not want
to gain e-,
therefore
low E.A.

19. Question: Arrange the following atoms in order of
increasing E. A.(lowest to highest):
F, Na, N, Rb, Al?
Explain your reasoning.

20. Answer: Rb < Na < Al < N < F metals: low I.E.
non – metals: high I.E.
Rb < Na < Al < N < F
5 shells shells shells shells shells
37 p p p p p +
more shielding  weaker less shielding  weaker “pull”
“pull” on valence shell e on valence shell e-
higher nuclear
charge  stronger
pull on e-

21. Electronegativity (E.N.)
Measure of an atom’s ability to attract a shared pair of electrons in a chemical bond
The difference in EN
determines the type
of bond that forms
between atoms

22. Atomic Radius I. E. E. A. E. N. decreases increases increases

23. Question: Which element has the highest electronegativity?
Which element has the lowest electronegativity?

24. Answer: MOST “ELECTRONEGATIVE” ELEMENTS TOP RIGHT HAND CORNER
MOST “ELECTROPOSITIVE”
ELEMENTS
BOTTOM LEFT HAND
CORNER

25. Reactivity Metals and non-metals generally react differently.
Metals react by _________________ electrons.
Metals have __________ E.N. and ______ I.E.
Non-metals react by ________________ electrons.
Non-metals have _______ E.N. and ______ E. A.
losing
low
low
gaining
high
high

26. Question: Which metal is the most reactive?
Which non-metal is the most reactive?
Which elements in the Periodic Table are the
least reactive?

27. MOST “REACTIVE” NON-METAL
Answer:
reactivity of metal decreases
reactivity of non-metal increases
MOST “ELECTRONEGATIVE”
NON-METAL
MOST “REACTIVE” NON-METAL
reactivity
of metal
increases
reactivity
of non-metal
decreases
MOST “ELECTROPOSITIVE”
METAL
MOST “REACTIVE” METAL

28. Question: Order the following metals in order of
increasing reactivity (lowest to highest)?
Cs Fe K

29. Answer: Fe < K < Cs Cs: more shells weaker hold on e-
lose e- easily, reactive
Fe: More protons, stronger hold on e-
Loses e- less easily
less reactive

30. Electronegativity (E.N.)
How powerful an atom is at “pulling” or “attracting” electrons to itself when it is bonded to another atom
The difference in EN
determines the type
of bond that forms
between atoms
Ionic
Covalent

31. Bonding Continuum Covalent Ionic  E. N. < 1.7  E. N. > 1.7
 EN = EN of atom 1 – EN of atom 2
1.7
Covalent
 E. N. < 1.7
Ionic
 E. N. > 1.7

32. Bonding Continuum – “Mostly” Covalent Bonds
C - H
C - F C F H
EXAMPLE:
 EN = EN of atom 1 – EN of atom 2
 EN C – H bond = 2.1 – = 0.4, therefore COVALENT
 EN C – F bond = 2.5 – = 1.5, therefore COVALENT

33. Bonding Continuum – “Mostly” Covalent Bonds
C - H
C - F
E. N. = 0
pure
covalent
Covalent
 E. N. < 1.7
Ionic
 E. N. > 1.7
0.5 <  E. N. < 1.7
 E. N. < 0.5
NON – POLAR COVALENT
e- shared “fairly equally”
POLAR COVALENT
e- shared “unequally”

34. Percent Ionic Character
% Ionic Character = ( E. N. )
3.3
Bond % Ionic Character
Na – Cl
2. C – H
3. C – F
(2.1/3.3) * 100 = 64 %
(0.4/3.3) * 100 = 12 %
(1.5/3.3) * 100 = 45 %