1. Periodic Table & Trends

2. History of the Periodic Table
• 1871 – Dimitri Mendeleev was the first
scientist to published an organized periodic
table. He arranged the elements according
to: 1. Increasing atomic mass 2. Elements
w/ similar properties were put in the same
row
• 1913 – Moseley arranged the elements
according to: 1. Increasing atomic number
2. Elements w/ similar properties were put
in the same column

3. The Periodic Law
• Mendeleev understood the ‘Periodic Law’
which states:
• The properties of the elements are periodic
function of their atomic number.

4. The Periodic Law
• Atoms with similar properties appear in
groups or families (vertical columns) on the
periodic table.
• They are similar because they all have the
same number of valence (outer shell)
electrons, which governs their chemical
behavior.

5. Period
Group
Metals
Non metals
Noble Gases
Representative

6. Group Names
Alkali
+1
Alkaline
Earth
Metals
+2 +3 -3 -2
Halogen
-1
Noble
Gases
0
H
1
He
2
Li
3
Be
4
B
5
C
6
N
7
O
8
F
9
Ne
10
Na
11
Mg
12
Al
13
Si
14
P
15
S
16
Cl
17
Ar
18

7. S & P block – Representative Elements
Metalloids (Semimetals, Semiconductors) – B,Si, Ge,
As, Sb, Te (properties of both metals &
nonmetals)
Columns – groups or families Rows - periods
METALS
TRANSITION METALS
NONMETALS

8. Periodic Groups
• Elements in the same column have similar
chemical and physical properties
• These similarities are observed because
elements in a column have similar e-
configurations (same amount of electrons in
outermost shell)

9. Periodic Trends
• Periodic Trends – patterns (don’t always
hold true) can be seen with our current
arrangement of the elements (Moseley)
• Trends we’ll be looking at:
1. Atomic Size and Radius
2. Ionization Energy
3. Electronegativity
4. Electron Affinity
5. Metallic Property

10. Atomic Size
• Size goes UP on going down a
group.
• Because electrons are added
farther from the nucleus, there is
less attraction.
• Size goes DOWN on going
across a period.

11. Atomic Radius
• Atomic Radius –
size of an atom
(distance from
nucleus to
outermost e-)

12. Atomic Radius Trend
• Group Trend – As you go down a column,
atomic radius increases
As you go down, e- are filled into orbitals that
are farther away from the nucleus (attraction
not as strong)
• Periodic Trend – As you go across a period
(L to R), atomic radius decreases
As you go L to R, e- are put into the same
orbital, but more p+ and e- total (more
attraction = smaller size)

13. Ionic Radius
• Ionic Radius –
size of an atom
when it is an
ion

14. Ionic Radius Trend
• Group Trend – As you go down a column, ionic
radius increases
• Periodic Trend – As you go across a period (L to
R), cation radius decreases,
anion radius decreases, too.
As you go L to R, cations have more attraction
(smaller size because more p+ than e-). The anions
have a larger size than the cations, but also
decrease L to R because of less attraction (more e-
than p+)

15. Ionic Radius

16. Ionic Radius
How do I remember this?????
The more electrons that are lost, the greater the
reduction in size.
Li+1 Be+2
protons 3 protons 4
electrons 2 electrons 2
Which ion is smaller?

17. Ionic Radius
How do I remember this???
The more electrons that are gained, the greater the
increase in size.
P-3 S-2
protons 15 protons 16
electrons 18 electrons 18
Which ion is smaller?

18. Ionization Energy
See Screen 8.12
IE = energy required to remove an electron from an
atom in the gas phase.
Mg (g) + 738 kJ ---> Mg+ (g) + e-

19. Ionization Energy
• Group Trend – As you go down a column,
ionization energy decreases
As you go down, atomic size is increasing (less
attraction), so easier to remove an e-
• Periodic Trend – As you go across a period (L to
R), ionization energy increases
As you go L to R, atomic size is decreasing (more
attraction), so more difficult to remove an e-
(also, metals want to lose e-, but nonmetals do
not)

20. Electronegativity
• Electronegativity-
tendency of an
atom to attract e-

21. Electronegativity Trend
• Group Trend – As you go down a column,
electronegativity decreases
As you go down, atomic size is increasing, so less
attraction to its own e- and other atom’s e-
• Periodic Trend – As you go across a period (L to
R), electronegativity increases
As you go L to R, atomic size is decreasing, so there
is more attraction to its own e- and other atom’s e-

22. Electron Affinity
A few elements GAIN electrons to form
anions.
Electron affinity is the energy change
when an electron is added:
A(g) + e- ---> A-(g) E.A. = ∆E

23. Electron Affinity of Oxygen
∆E is EXOthermic
because O has an
affinity for an e-.
[He] 
 
  
O atom
EA = - 141 kJ
+ electron
O [He] 
 
  
- ion

24. Electron Affinity of Nitrogen
∆E is zero for N- due
to electron-
electron
repulsions.
EA = 0 kJ
[He] 
  
N atom 
[He] 
  
N- ion 
+ electron

25. Reactivity
• Reactivity – tendency of an atom to react
• Metals – lose e- when they react, so metals’
reactivity is based on lowest Ionization Energy
(bottom/left corner) Low I.E = High Reactivity
• Nonmetals – gain e- when they react, so
nonmetals’ reactivity is based on high
electronegativity (upper/right corner)
High electronegativity = High reactivity

26. Metallic Character
• Properties of a Metal – 1. Easy to shape
2. Conduct electricity 3. Shiny
• Group Trend – As you go down a column, metallic
character increases
• Periodic Trend – As you go across a period (L to
R), metallic character decreases (L to R, you are
going from metals to non-metals

27. Effective Nuclear Charge,
Z*
• Z* is the nuclear charge experienced by
the outermost electrons.
• Explains why E(2s) < E(2p)
• Z* increases across a period owing to
incomplete shielding by inner electrons.
• Estimate Z* by --> [ Z - (no. inner electrons) ]
• Charge felt by 2s e- in Li Z* = 3 - 2 = 1
• Be Z* = 4 - 2 = 2
• B Z* = 5 - 2 = 3 and so on!

28. Effective Nuclear Charge, Z*
• Atom Z* Experienced by Electrons in
Valence Orbitals
• Li +1.28
• Be -------
• B +2.58
• C +3.22
• N +3.85
• O +4.49
• F +5.13
Increase in Z*
across a period

29. Shielding
• The less attracted to the nucleus, the more
shielded, thus lesser effective nuclear
charge.
• The effective nuclear charge on those outer
electrons is less, and so the outer electrons
are less tightly held.

30. Effective Nuclear Charge
• What keeps electrons from simply flying off
into space?
• Effective nuclear charge is the pull that an
electron “feels” from the nucleus.
• The closer an electron is to the nucleus, the
more pull it feels.
• As effective nuclear charge increases, the
electron cloud is pulled in tighter.

31. General Periodic Trends
• Atomic and ionic size
• Ionization energy
• Electron affinity
• Electronegativity
Higher effective nuclear charge.
Electrons held more tightly
Smaller orbitals.
Electrons held more
tightly.