periodictrends- chemistry for eight graders

This is a PPT about periodic trends

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. Perio
d
GroupMetalsNon metalsNoble GasesRepresentative

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

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

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

13. 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)

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

15. 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
+
)

16. Ionic Radius

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

18. Ionic Radius
How do I remember this???
The more electrons that are gained, the greater the
increase in size.
P
-3
S
-2
protonsprotons1515 protons 16protons 16
electrons 18electrons 18 electrons 18electrons 18
Which ion is smaller?Which ion is smaller?

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

20. 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)

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

22. 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
-

23. Electron AffinityElectron Affinity
A few elements A few elements GAINGAIN electrons to form electrons to form
anionsanions..
Electron affinity is the energy change Electron affinity is the energy change
when an electron is added:when an electron is added:
A(g) + e- ---> AA(g) + e- ---> A
--
(g) E.A. = ∆E(g) E.A. = ∆E

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

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

26. 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

27. 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

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

29. EffectiveEffective Nuclear Charge, Z* 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* Increase in Z*
across a periodacross a period

30. 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.

31. 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.

32. General Periodic TrendsGeneral Periodic Trends
•Atomic and ionic sizeAtomic and ionic size
•Ionization energyIonization energy
•Electron affinityElectron affinity
•ElectronegativityElectronegativity
Higher effective nuclear charge.
Electrons held more tightly
Smaller orbitals.
Electrons held more
tightly.